/ 

SOLDBILIT 



ER AT 18°. 



ci 



Br 



NO, 



32.95 
3.9 



137.5 
6.0 



92.56 
12.4 



30.34 
2.6 



Na 



5.42 



88.76 




Mg 



55.81 
5.1 



177.9 
8.1 



03.1 
4.6 



fibre ATP 7. { 



18.2 
4.1 



4.44 
1.06 



Book 



'- 



_ 



0.0087 
0.0 2 14 



83.97 
7.4 



°i \ $ 



74.31 
4.0 



126.4 
4.7 




Zn 



203.9 
9.2 



478.2 



419 
6.9 



0.005 

0.0 S 5 



117.8 
4.7 



183.9 
5.3 



Pb 



1.49 
0.05 



0.598 
0.02 



0.08 
0.0-2 



0.06 
0.002 



51.66 
1.4 



150.6 
3.16 



The upper number in each square gives the number of grams of the 
anhydrous salt held in solution by 100 c.c. of water. The lower number is 
the molar solubility, i.e., the number of moles contained in one liter of the 
saturated solution. The numbers for small solubilities have been abbreviated. 
Thus 0.0 tt 4 -» 0.0000004. For some other solubilities, see page 131. 



Bs tbe same autbor 

A LABORATORY OUTLINE OF 

GENERAL CHEMISTRY 

NEW EDITION, 1908 

Fourth Edition, Revised 

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of the University of Michigan 

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New York, THE CENTURY CO. 
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GENERAL CHEMISTET 
FOR COLLEGES 



BY 

ALEXANDER SMITH 

Professor of Chemistry, and Head of the Department, 
Columbia University 



SECOND EDITION 

ENTIRELY REWRITTEN 




OXE HUNDRED AXD FOURTH 
THOUSAND 



NEW YORK 

THE CENTURY CO, 
1918 






QH3I 



Copyright, 1905, 1906, 1908, 1916, 

BY 

THE CENTURY CO. 



First Edition, May, 1908 

Reprinted October, 1908 ; April, 1909 ; May, 1910; May, 1911; 

April, 1912 ; April, 1913 ; May, 1914 ; January, 1915 

Second Edition, January, 1916 

Reprinted July, 1916 ; August, 1916; 

September, 1916 ; October, 1916 

June, 1917; October, 1917; May, 1918 

Reprinted, September, 1918 

'Rpprfotad. Oo.tober. 1918 

Reprinted, November, 1918 



Library of Congreaa 
By transfer from 
War Department, 



RrtMri 29 



PREFACE TO THE FIRST EDITION 



The present work differs from the Author's "Introduction to 
General Inorganic Chemistry" in being intended for pupils who 
can devote less time to the study of the science, and whose needs 
can be satisfied by a less extensive course. It resembles the 
larger work in the arrangement of the contents and in the general 
method of treatment. The matter, and particularly the theoreti- 
cal matter, however, has been simplified and has been confined 
strictly to the most fundamental topics. Such parts of the theory 
as are thus given, are presented with the same fullness as before, 
and are illustrated and applied with all the persistence needed to 
insure full apprehension and, ultimately, spontaneous employment 
by the student. Such parts as could not be treated in this way, 
within the limits set by the plan of the book, have been omitted. 
Methods materially different from those used in the " Introduction" 
have been employed in presenting many topics. Conspicuous 
differences of this kind will be noted particularly in the treatment 
of combining proportions, formulae and equations, molecular and 
atomic weights, chemical equilibrium, ionic substances and their 
interactions, and the theory of precipitation. 

The writer desires to express his profound gratitude to the many 
chemists who have made valuable criticisms and suggestions. 
Most of these comments applied to the "Introduction to General 
Inorganic Chemistry," but many of them have been used in 
preparing this work (General Chemistry for Colleges), and all will 
be considered in the second edition of the larger book. 

For critical reading of the whole of the proofs of the present 
work, the writer desires especially to thank Messrs. A. T. McLeod 
and Alan W. C. Menzies of the University of Chicago. Other cor- 
rections and suggestions will be gladly received by the author. 

Alexander Smith. 
Chicago, April, 1908. 



I 



PREFACE TO THE SECOND EDITION 



In preparing the second edition, the entire book has been re- 
written. The introduction to the subject has been improved and 
greatly simplified, and several difficult topics have been trans- 
ferred to later chapters. The explanations of the theoretical 
subjects, and of the methods of making calculations have been 
clarified and additional illustrations have been given. In view 
of its importance to prospective students of biology and medicine, 
osmotic pressure is treated in greater detail. To add greater 
interest to the study of the science, and because of their edu- 
cational value, the historical references have been expanded, 
many more applications of chemistry have been discussed, and 
the number of figures has been considerably increased. Exten- 
sive new sections on oxidation and reduction, and on various 
methods of writing equations, on radio-activity, and on electro- 
motive chemistry have been added. Briefer new sections on 
atomic numbers, colloids, foods, explosives, water treatment, and 
many other subjects, have been included. New pedagogical de- 
vices have been introduced. So far as recent advances can be 
apprehended and applied by first-year college students, the treat- 
ment has been brought up to date. 

The author is greatly indebted to Messrs. P. C. Haeseler, and 
Herbert E. Eastlack for much assistance in reading the proof and 
for many valuable suggestions. 

Alexander Smith. 

New York, January, 1916. 



vu 



CONTENTS 



Chapter Page 

I. The Chemical View of Matter 1 

II. Chemical Change and the Methods of Studying It. . . 11 

III. Oxygen 25 

IV. Atomic Weights, Symbols, Formula, and Equations ... 40 
V. Hydrogen 49 

VI. Valence. Calculations 61 

VII. The Measurement of Quantity in Gases. Relations 

Between Structure and Behavior of Matter 70 

VIII. Water 85 

•IX. Molecular Weights and Atomic Weights 100 

X. Solution 121 

XI. Hydrogen Chloride. Calculations 141 

XII. Chlorine 154 

XIII. Energy and Chemical Change 167 

XIV. Chemical Equilibrium 177 

XV. The Halogen Family 192 

XVI. Dissociation in Solution 210 

XVII. Ozone and Hydrogen Peroxide 219 

XVIII. Ionization 226 

XIX. Ionic Substances and their Interactions 245 

XX. Sulphur and Hydrogen Sulphide 264 

XXI. The Oxides and Oxygen Acids of Sulphur 275 

XXII. Selenium and Tellurium. The Classification of the 

Elements 293 

XXIII. Oxides and Oxygen Acids of the Halogens. Oxida- 

tion and Reduction 306 

XXIV. The Atmosphere. The Helium Family 328 

XXV. Nitrogen and Ammonia 338 

XXVI. Oxides and Oxygen Acids of Nitrogen 347 

XXVII. Phosphorus 362 

ix 



X CONTENTS 

Chapter Page 

XXVIII. Carbon and the Oxides of Carbon 375 

XXIX. The Hydrocarbons. Flame 389 

XXX. The Carbohydrates and Related Substances 402 

XXXI. Organic Acids and Salts. Alcohols, Esters. Foods.. 412 

XXXII. Silicon and Boron 425 

XXXIII. The Base-forming Elements 434 

XXXIV. The Metallic Elements of the Alkalies: Potassium 

and Ammonium 443 

XXXV. Sodium and Lithium. Ionic Equilibrium Considered 

Quantitatively 457 

XXXVI. The Metallic Elements of the Alkaline Earths . . . 473 

XXXVII. Copper, Silver, Gold 500 

XXXVIII. Glucinum, Magnesium, Zinc, Cadmium, Mercury. The 

Recognition of Cations in Qualitative Analysis . . . 523 

XXXIX. Electromotive Chemistry 539 

XL. Aluminium and the Metals of the Earths 553 

XLI. Germanium, Tin, Lead 567 

XLII. Arsenic, Antimony, Bismuth 582 

XLIII. The Chromium Family. Radium 595 

XLIV. Manganese 617 

XLV. Iron, Cobalt, Nickel 625 

XLVI. The Platinum Metals 644 

APPENDIX 648 



GENERAL CHEMISTRY FOR COLLEGES 



GENERAL CHEMISTRY FOR COLLEGES 



CHAPTER I 

THE CHEMICAL VIEW OF MATTER 

Chemistry is a science which deals with all forms of matter. It 
considers the natural kinds, such as rocks and minerals, as well as 
materials like fat and flour obtained from animals or plants. It 
deals also with artificial products like paints or explosives. When 
we wish information about any specimen or kind of matter, we 
consult a chemist. Now chemists have worked out a point of 
view which enables them to attack any problem connected with 
matter in a systematic manner and to state the results in a clear 
and simple way. To learn something of chemistry, we must 
acquire this point of view and master the technical language the 
chemist uses in stating and discussing his results. 

Properties. — Suppose that a piece of rusty iron is submitted 
to the chemist. He at once -examines the rust and notes that it 
is solid, reddish-brown in color and earthy in appearance. He 
separates some of it from the iron and finds it to be brittle, that is, 
easily broken and capable of being pulverized in a mortar. He 
finds that its density is about 4.5, that is to say, 1 c.c. (Appendix I) 
of it weighs about 4.5 g. On heating some of it in a flame, he finds 
that it does not melt, and must therefore have a very high melting- 
point. .These qualities he calls properties, and more especially 
physical properties. Since all specimens of iron-rust show exactly 
the same properties, he often calls them specific physical proper- 
ties, because they are properties shown by all specimens of a par- 
ticular species of matter. 

After removing any rust by filing or scraping, the chemist ex- 
amines the iron, and finds a fresh, clean surface to be almost white 
and metallic in appearance. The metal is tenacious, so that it can 

1 



Z COLLEGE CHEMISTRY 

be bent but not easily broken. It is ductile and can therefore be 
drawn out into wire. He finds that its density is about 7.5, and 
that the metal is incapable of being melted in an ordinary flame. 
In addition, he finds it to be strongly attracted by a magnet, 
while rust is not attracted. 

The chemist, then, studies what he calls the specific physical 
properties of each material, in order that he may be able to recog- 
nize various materials. 

Substances. — All specimens of iron show one set of proper- 
ties and all specimens of iron-rust show a different set, peculiar to 
rust. The chemist calls any definite variety of matter, all speci- 
mens of which show the same properties, a substance. Iron is one 
substance and rust another. A substance is recognized by its 
properties. 

The point of view of the chemist thus consists in describing any 
material by ascertaining whether it is made up of one, or of more 
than one substance. He describes it by naming the substances 
which, by a study of their properties, he has found in it. 



Two Illustrations of the Study and Description of Ma- 
terials. — If a piece of granite is examined by a chemist, he 
observes at once that it is spotted in appearance, and made up of 
several crystalline materials of differing nature. He therefore 
breaks it up and studies the properties of the fragments. Some of 
the fragments of granite are dark and with a penknife can easily 



a 



(5 



Fig. 1. 



Fig. 2. 




Fig. 3. 



be split into transparent sheets, thinner than paper. These par- 
ticular fragments are in all respects like mica (Fig. 1). This sub- 
stance is a mineral which, in certain neighborhoods, occurs in 
large masses, and sheets of it ("isinglass") are used to fill the 
openings in stoves. Others of the fragments are clear like glass, 
and are very hard (see Appendix II), and have all the properties 



THE CHEMICAL VIEW OF MATTER 6 

of quartz or rock crystal (Fig. 2), which is another substance 
well known to the chemist. The remaining fragments are less 
clear than is quartz, and are not so hard. They can be split into 
layers, but not nearly so easily as can mica. They form oblong 
crystals, differing in this also from quartz, which shows hexagonal 
crystals.* This substance is felspar (Fig. 3). Thus the chemist 
studies the physical properties of the fragments, and finds that 
there are three different substances in granite. He reports that 
the components of granite are mica, quartz, and felspar. 

When flour is examined by the chemist, it appears to the eye to 
be all alike. Under the microscope, even, all he can learn is that 
it consists largely of grains, which have the 
characteristic appearance (first property) of 
grains of starch (see Fig. 108, p. 403). He 
places some flour on a square piece of cheese- 
cloth and encloses it by tying with a thread 
(Fig. 4). On kneading the little bag in a vessel 
of water, the water becomes milky. When the 
milky water stands, the white material settles 
to the bottom, the water can be poured off, 
and the deposit can be dried. This white sub- FlG ' 4 ' 

stance, when boiled with water, gives an almost clear liquid which 
jellies on cooling. This is another property of starch. A little 
tincture of iodine (solution of iodine in alcohol), dropped on a 
part of the starch, causes the latter to turn blue. This is a very 
characteristic property of (and therefore test for) starch. When 
the bag of flour is kneaded persistently in water which is frequently 
changed, the material finally ceases to render the water milky. 
The starch has all been washed out. When the bag is now opened, 
a sticky material is found in it. This is called gluten. The chemist 
therefore finds that the flour contains starch and gluten. He 
learns this by separating the components. 

Law of Component Substances, — Every material can be 
described as being composed of one substance, or as being a mixture 

* Crystals (see also Index) are natural forms, of geometrical outline, which 
solid substances assume. Usually each substance has a more or less distinct 
form of its own, the particular angles at which the faces meet being peculiar to 
the substance. Its individual crystalline form is therefore a specific physical 
property of each substance. 




4 COLLEGE CHEMISTRY 

of two or more component substances, each of which has a definite 
set of specific physical properties. This is the first and most 
fundamental law of chemistry. This conception was first clearly 
stated by Lomonossov (1742), a Russian author, statesman, and 
.chemist (1711-1765). 

Mixtures and Impurities. — A material containing more 
than one component substance is called a mixture. The charac- 
teristic of a mixture is that each of the component substances, 
although mixed with the others, possesses exactly the same prop- 
erties as if it were present alone. No one of the components 
affects any other component, or alters any of its properties. 
Granite and flour are typical mixtures. 

When a specimen is composed mainly of one substance, and 
contains only minute amounts of one or more other substances, 
it is frequently spoken of as a specimen of the main substance 
containing certain specified substances as impurities. To be called 
an impurity, the foreign matter need not be dirty or offensive. 
Thus,; Qpmmon salt usually contains a little magnesium chloride, 
a white crystalline solid, as an impurity, and it is this impurity 
which becomes damp in wet weather. Again, compounds of 
lime and magnesium are common impurities in drinking water. 

Components. — The ingredients of a mixture are called the 
components (Lat., put with), because they are simply placed to- 
gether, without change, and can be separated without change. 

Bodies or Specimens. — It will be seen that substance is a 
general term, like the word "dog," covering the whole species. 
The substance iron includes all the iron in the universe. When 
we refer to a particular piece of iron, we call it a body or a speci- 
men. If the body is homogeneous (all parts alike), it may be 
made of a single substance. If it is heterogeneous (differing in 
different parts) it is a specimen of a mixture like granite. 

The Rusting of Metals. — If we return once more to the 
subject of rusty iron, we find another point which interests the 
chemist. If the iron is kept moist — for example, by lying in 
the grass or partly immersed in water — the layer of rust gradu- 



THE CHEMICAL VIEW OF MATTER O 

ally becomes thicker, and the core of iron becomes thinner, until 
it finally disappears. The rust seems to be formed from the iron, 
in presence of air and moisture. The iron, particle by particle, 
loses the properties of iron and simultaneously acquires those of 
rust. Now the chemist is concerned, not only with recognizing* 
substances, but also with the ways in which substances change and 
new substances are produced. 

Several other familiar metals rust, as does iron, but the change 
is slower. Thus, lead rusts (tarnishes) slowly, and zinc still more 
slowly. The change can be hastened by 
heating. For example, if some lead is 
melted in a porcelain crucible (Fig. 5) 
and is stirred with an iron wire, a dirty 
yellow powder collects on the surface. 
Gradually more and more of the powder 
is formed and less and less of the metallic 
lead remains, until at last all the metal is 
gone. Melted tin, when treated in the 
same way, gives a white powder. 

Explanation of Rusting. — The first 
fact which seemed to throw light on the sub- 
ject was discovered by a French physician, FlG - 5 - 
Jean Rey (1630), who found that the rusts of tin and lead, made 
by heating and stirring, were heavier than the original pieces of 
metal. He inferred, correctly, that the additional material which 
caused the increase in weight came from the air. He imagined, 
however, that the rust was not a new substance, but a sort of 
froth, and therefore a mixture of air with the metal. Other in- 
vestigators, such as Hooke (1635-1703) and particularly Mayow 
(1645-1679), in England, explained the increase in weight by sup- 
posing that some material from the air had combined with the 
metal. In other words, iron, for example, was one substance 
composed of iron only, and rust was another substance, made by 
union of iron and a material from the air, and not a mere mixture. 

It was Lomonossov (1756) who first proved by an experiment 
that the extra material did come from the air. He placed some 
tin in a flask, sealed up the mouth of the vessel, and weighed the 
whole. The flask was then heated and the tin was converted into 




6 



COLLEGE CHEMISTRY 



the white powder. So long as the flask remained sealed, no change 
in weight was found to have occurred. When the mouth of the 
flask was opened, however, some air rushed in, and the total 
weight was then found to be greater. Evidently, during the heat- 
ing, a portion of the original air had forsaken the gaseous condi- 
tion and joined itself to the tin to form the powder. This left a 
partial vacuum in the flask, and more air entered when the latter 
was opened. Eighteen years later the same experiment was made 
by Lavoisier, who drew the same conclusion. The rusting of 
other metals was found to be due to the same cause. Lavoisier 
named the gas, taken from the air, oxygen. 

The conclusion can be confirmed in various ways. For example, 
when the air is pumped out of the flask before it is sealed, the metal 
can be heated in the vacuum indefinitely without rusting. 



■. : 
.;; 

;■■: 
- 



Experiment to show the Nature of Rusting. — That a 
part of the air is consumed when iron rusts is easily proved. We 
moisten the interior of a test-tube and sprinkle some powdered 
iron so that it covers and adheres to the 
whole interior surface. We then set the 
tube mouth downwards in a dish of water 
(Fig. 6) . At first, the pressure of the water 
compresses the air in the tube very slightly, 
and the water ascends above the mouth to 
the extent of a small fraction of an inch 
only. As the moist iron slowly rusts, how- 
ever, the oxygen is gradually removed, and 
the pressure of the atmosphere outside slowly 
pushes the water farther up the tube. After an hour or more, the 
water has ascended about one-fifth of the total distance towards 
the top of the tube. Evidently part of the air has forsaken the 
gaseous condition, and the water has been forced up to take its 
place. Inspection now shows some reddish particles, where rusting 
has taken place. The rust, then, is made up of a part of the iron 
and all of the oxygen that the tube contained. 

Of course, much of the iron powder is still gray, and has not 
rusted. The air in the tube did not contain oxygen enough to 
combine with all the iron. The iron that remains is as little able 
to rust in the remaining gas as in a vacuum. 



Fig. 



THE CHEMICAL VIEW OF MATTEB 7 

Incidentally we learn from this experiment that atmospheric 
air contains about one-fifth (20 per cent) oxygen by volume. 
The remaining four-fifths is almost all nitrogen (79 per cent), a 
substance which combines with very few materials, while the 
balance (1 per cent) is made up of gases which do not enter into 
combination with any known substance. If lead, tin, or zinc 
had been heated in an enclosed volume of air, they likewise would 
have taken out the 20 per cent of oxygen and would have left 
the other gases. 

The Law of Chemical Change. — The three examples of 
rusting show that specimens of matter can lose their original 
properties and acquire new ones. Since a substance is "a spe- 
cies of matter, with a constant set of properties," we are compelled 
to decide that, when a material changes its properties, it has, in 
doing so, become a new substance. This consideration calls to 
our attention the second of the fundamental laws of chemistry, 
namely, that the material forming one or more substances (such as 
oxygen and iron), without ceasing to exist, may be changed into 
one or more entirely different substances. Such a change is called 
a chemical change, or action, or interaction, or reaction. 

The commoner kinds of chemical actions can be divided, for 
convenience, into four varieties. We can now define the first of 
these. 

First Variety of Chemical Change: Combination. — In 

each case of rusting, two substances (a gas and a metal) came 
together to form a third substance (an earthy powder). Appar- 
ently two substances may come together in two different ways. 
They may form a mixture, in which both substances are present 
and retain their properties, or they may come together to form a 
single substance with different properties. When two (or more) 
substances unite to form one substance, the change is called chem- 
ical combination or union. The product is called a compound 
substance. 

We are very careful never to speak of a compound substance 
as a mixture. Rust is not a mixture of iron and oxygen; it shows 
none of the properties of either. Nor do we call a mixture (like 
granite) a compound, or the operation of mixing, combination or 



8 COLLEGE CHEMISTRY 



.. 



union. These are technical words, in chemistry and, to avoid 
confusion, may be used only with due regard to their technical 
meanings. 

Constituents. — As we have seen, we speak of the substances 
in a mixture as the components. When we wish to refer to the 
forms of matter which are chemically united in a compound, we 
call them the constituents (Lat., standing together) of the compound 
substance. Thus, iron and oxygen are the constituents of rust. 

The chemist separates (p. 3) the components of a mixture, for 
that is all that is necessary. He liberates the constituents of a 
compound, however, because they are bound together in chemical 
combination. 

The names given to compounds are usually devised so as to 
indicate the nature of the constituents. Thus, iron-rust is oxide 
of iron (or ferric oxide, from Lat. ferrum, iron). The yellowish 
powder from lead is lead oxide or oxide of lead, and the white 
powder from tin is oxide of tin. 

A Condensed Form of Statement. — We may represent a 
chemical combination, or indeed any kind of chemical change, in a 
condensed form, thus: 

Iron + Oxygen — > Oxide of iron (ferric oxide). 

Each name stands for a substance. Two substances in contact 
with one another (mixed), but not united chemically, are con- 
nected by the -f- sign. The arrow shows where the chemical 
change comes in, and the direction of the change. We read the 
statement thus : Iron and oxygen brought together under suitable 
conditions undergo chemical change into oxide of iron, called also 
ferric oxide. Similarly we may write: 

Lead + Oxygen — » Oxide of lead. 
Tin + Oxygen — » Oxide of tin. 

The Increase in Weight in Rusting. — As we have seen, 
the process of rusting is accompanied by a slow increase in the 
weight of the solid, due to the gradual addition of oxygen to the 
metal. Now, this increase in weight ceases of its own accord, 
when a certain maximum has been reached. This occurs when 



THE CHEMICAL VIEW OF MATTER 9 

the last particles of the metal have disappeared. Thus, the lead 
gains in weight until every 100 parts of the metal have gained 7.72 
parts of oxygen, and the tin until every 100 parts have gained 
26.9 parts of oxygen. When these increases have occurred, the 
metal is found to have been all used up, and prolonged heating 
and stirring cause no further union with oxygen and no further 
change in weight. This fact, that each substance limits itself 
of its own accord to combining with a fixed proportion of the other 
substance, in forming a given compound, is one of the most strik- 
ing facts about chemical combination. In mixtures, any propor- 
tions chosen by the experimenter may be used. In chemical 
union, the experimenter has no choice; the proportions are de- 
termined by the substances themselves. Thus, 100 parts of iron 
when turning into ordinary red rust take up 43 parts of oxygen, 
no more and no less. 

This fact enables us to make our condensed statements more 
specific and complete by including in them the proportions by 
weight used in the chemical change: 

Iron (100) + Oxygen (43) -> Ferric oxide (143). 
Lead (100) + Oxygen (7.72) -> Oxide of lead (107.72). 

The following numbers, which represent the same proportions 
by weight, are the ones commonly used by chemists: 

Iron (111.68) + Oxygen (48) ^Ferric oxide (159.68). 

Summary. — Thus far, we have learned that chemistry deals 
with substances and their physical properties, and with the changes 
which substances undergo. We have discussed and defined a 
number of important words expressing fundamental chemical 
ideas. Finally, we have touched upon the weights of the ma- 
terials used in chemical change, a subject of great importance 
which will be more fully developed in a later chapter. 

Exercises. — 1. Take one by one the words or phrases printed 
in black type and the titles of the sections in this chapter, and 
endeavor to recollect what you have read about each. In each 
case try, (a) to recall the meaning and to state it in your own 
words; (b) to recall the facts associated with, and the reasoning 
which lead up to the point in question; (c) to recall examples 



10 COLLEGE CHEMISTRY 

illustrating the conception and to apply the conception in detail 
to each example. Whenever memory fails to give a perfectly 
clear report of the matter in hand, the text must be read and 
re-read until the essential point can be repeated from memory. 

Use the same method in all future chapters. A useful prac- 
tice is to employ a pencil as you read and to underline systemati- 
cally all the important facts and statements, and then to go back 
and apply to each marked place the process described above. 

2. Define the following terms: Specific gravity, tenacity, melt- 
ing-point, specific physical property, pure body, vacuum. 

3. Is it logical to say "pure substance?" 

4. Why do we decide that granite is a mixture and iron a 
single substance? 

5. Do the statements in the text indicate that air is a mixture 
or a compound? 

6. What weight of oxygen would be required to convert 25 
grams of lead into oxide of lead? 

7. Make a list of the technical words we have defined, and place 
the definition opposite to each. 

8. What weight of tin would be contained in 15 grams of oxide 
of tin? 

9. If any of the following are mixtures, mention the facts which 
show them to contain more than one substance: (a) muddy water, 
(6) an egg, (c) milk. 

10. In recognizing a specimen to be quartz, does the chemist 
consider (a) the weight, (6) the temperature, (c) the length of the 
specimen? If not, why not? 

11. Give a list of the specific properties mentioned in this 
chapter. 



CHAPTER II 

CHEMICAL CHANGE AND THE METHODS OF 
STUDYING IT 

We must now take up two new examples of chemical change. 
They will aid us in introducing one or two additional conceptions 
and laws. These are continually used by the chemist, and without 
them we cannot begin the systematic study of the science. 

Another Case of Combination: Iron and Sulphur. — 

Since oxygen is an invisible gas, there is a slight difficulty in real- 
izing that rusting consists in the union of two substances — this 
gas and a metal. The present example is less interesting histori- 
cally, but it is simpler because both substances are visible and are 
easily handled. The case of iron and sulphur will enable us to 
illustrate the same point of view and to practice the application 
of the same technical words. It will also introduce us to two 
manipulations — filtration and evaporation — which are fre- 
quently used by the chemist. 

We begin by observing the physical properties of the two sub- 
stances. Those of iron have already been noted (pp. 1-2).* Sul- 
phur is a pale-yellow substance of low specific gravity (sp. gr. 2). 
It is easily melted (m.-p. 112.8° C). It does not dissolve in water 
— that is, it does not mix completely with and disappear in water, 
as sugar does on stirring. It does dissolve readily in carbon 
disulphide, however. It crystallizes in rhombic forms (Fig. 7). 
It is not attracted by a magnet. 

* References to previous pages are used in order to save needless repetition 
in writing. The beginner requires endless repetition in his reading, however, 
and must form the habit of examining, in conjunction with the current text, the 
parts referred to. The passages cited are, by the reference, made part of the 
current text, which will usually not be clear without them. The same remark 
applies to topics referred to by name. Such topics must be sought in the index. 
All terms, and especially those borrowed from physics, if not perfectly 
familiar, must be looked up in a work on physics or in a dictionary. 

11 



12 



COLLEGE CHEMISTRY 




Study of the Mixture, before Combination. — Now, if 
some iron filings and pulverized sulphur are stirred together in a 
mortar, the result is a mixture. True, the color is not that of 
either substance, but with a lens particles 
of both substances can be seen. Passing a 
magnet over the mixture will easily remove 
a part of the iron, and with the help of a lens 
and a needle the mixture can be picked apart 
particle by particle, completely. We can sep- 
arate the components of the mixture more ex- 
M 7 * peditiously, however, by using manipulations 

based upon certain suitable properties. Thus, sulphur dissolves in 
carbon disulphide while iron does not. If, there- 
fore, a part of the mixture is placed in a dry test- 
tube along with some carbon disulphide (Fig. 8), 
and is shaken, the liquid dissolves the sulphur 
and leaves the iron. To complete the separa- 
tion, the iron must be removed from the liquid 
by filtration, and the sulphur recovered by evap- 
oration of the carbon disulphide. 

Filtration. — Iron, or any solid, when it is Fia 8 * 

mixed with a liquid or with a solution (like the solution of sulphur 

in carbon disulphide) is said to be sus- 
pended in the liquid. If the solid is one 
that settles rapidly, the liquid may be 
separated from the solid, in a rough 
way, by pouring off as much of the clear, 
supernatant liquid as possible. This is 
called decantation. 

A complete separation is effected by 

pouring the mixture on to a cone of 

filter paper supported in a glass funnel 

(Fig. 9). The liquid, together with 

anything that may be dissolved in it, 

fig. 9. run s through the pores of the paper and 

down the hollow stem of the funnel. The liquid is then called the 

filtrate. The particles of the suspended solid are too large to pass 

through the pores, and so collect on the surface of the filter paper. 






CHEMICAL CHANGE AND METHODS OF STUDYING IT 13 

This operation, like everything the chemist does, takes advantage 
of the physical properties of the various materials. 

The material remaining on the paper (the residue), when dry, 
is wholly attracted by a magnet and shows all the other properties 
of iron. 

Evaporation. — To recover the sulphur, the solution in carbon 
disulphide — the filtrate — is poured into a porcelain evaporat- 
ing dish (Inflammable! Keep flames away). When the vessel 
is set aside, the liquid gradually passes off in vapor (e-vapor-ates) . 
Sulphur, however, does not evaporate at room temperature and 
remains as a residue, in the form 
of crystals of rhombic outline in 
the bottom of the dish (Fig. 10). 
Here, again, physical properties 
have been utilized. 

Since the physical properties FlG * 10 * 

of two substances are not changed by mixing, we have thus used 
the properties of the iron and sulphur so as to separate them once 
more. The iron is on the paper; the sulphur is in the dish. 

Combination of Iron and Sulphur. — But iron and sulphur 
are capable of combining. If we alter the conditions by raising 
the temperature of some of the dry mixture, as we did in causing 
lead to rust rapidly, chemical union sets in. When we place some 
of the original mixture of iron and sulphur into a clean test-tube 
and warm it, we soon notice a rather violent development of heat 
taking place, the contents begin to glow, and what appears to be a 
form of combustion spreads through the mass. The heating em- 
ployed at the start falls far short of accounting for the much 
greater heat produced. When these phenomena have ceased, 
and the test-tube has been allowed to cool, we find that it now 
contains a somewhat porous-looking, black solid. This material 
is brittle; it is not magnetic; it does not dissolve in carbon di- 
sulphide; and close examination, even under a microscope, does 
not reveal the presence of different kinds of matter. This sub- 
stance is known to chemists as ferrous sulphide and, as we see, its 
properties are entirely different from those of its constituents. 

In this connection we must not omit to notice that, as in rusting, 



14 COLLEGE CHEMISTRY 



a certain fixed proportion will be used in forming the compound 
We find that, for 7 parts of iron, almost exactly 4 parts by weight 
of sulphur are required. If more iron is put into the original 
mixture, then some unused iron will be found in the mass after the 
action. If too much sulphur is employed, some may be driven off 
as vapor by the heat and any that remains, beyond the correct 
proportion, can be dissolved out of the ferrous sulphide with car- 
bon disulphide. The sulphur which has combined with the iron, 
however, is no longer present as sulphur — it has no longer the 
properties of sulphur, and therefore cannot be dissolved out : 
Iron (55.84) + Sulphur (32.07) -> Ferrous sulphide (87.91). 

Another Illustration: Mercuric Oxide. — It has long been 
known that air contains an active and an inactive gas. The 
Chinese called them yin and yang, respectively. Mayow (1643- 
1679) showed that the active gas caused rusting, that it was 
absorbed by paint (really by the linseed oil) in " drying," that it 
supported combustion of wood and sulphur, and that it is neces- 
sary to life, being absorbed by the blood from the air entering the 
lungs. It was not until 1774, however, that a pure specimen of 
this gas was obtained, by Priestley, and was rec- 
ognized to be a special kind of gas different from 
ordinary air. The gas (later to be named oxygen) 
was made by Priestley from mercuric oxide, a 
bright red, rather heavy powder. When the oxide 
is heated (Fig. 11), we find that a gas is given off. 
This gas is easily shown to be different from air, 
since a glowing splinter of wood is instantly re- 
Fig. n. lighted on being immersed in it. The gas is pure 

oxygen. During the heating, we notice also that a metallic coating 
appears on the sides of the tube, in the form of a sort of mirror. 
Apparently the vapor of some metal is coming off with the oxygen 
and condensing on the cool parts of the tube. As this shining sub- 
stance accumulates it takes the form of globules, which may be 
scraped together. It is, in fact, the metal mercury, or quicksilver. 
If the heating continues long enough, the whole of the red powder 
eventually disappears, and is converted into these two products. 

Second Variety of Chemical Change: Decomposition. — ■ 

Priestley's experiment introduces to us a second, and very common 



. 




CHEMICAL CHANGE AND METHODS OF STUDYING IT 15 

kind of chemical action. The first variety was combination or 
union (p. 7). The second is called decomposition. It consists in 
starting with a single substance (here mercuric oxide) and splitting 
it into two (or more) substances, which differ in properties from the 
substance taken and from one another. Here, the red powder 
gave mercury, a liquid metal, and oxygen, a colorless gas. 

Simple and Compound Substances. — We have seen that 
two (or more) substances, like lead and oxygen, can combine to 
form a compound substance. Are all substances, then, com- 
pounds? We find that some are not. We have never succeeded 
in obtaining lead, or oxygen, or iron, or tin, or sulphur by com- 
bining any two substances. We can decompose mercuric oxide 
by heat, and we have other ways of decomposing compounds like 
oxide of tin and ferrous sulphide, but we have never succeeded in 
decomposing the mercury or the oxygen, the iron or the sulphur 
themselves. Substances which we are not able, at will, to decompose 
into, or to make by chemical union from, other substances are called 
simple or elementary substances. The distinction between simple 
and compound substances was first drawn by Lomonossov in 1741. 
Later, and independently, it was stated very clearly by Lavoisier 
(1789). 

Several substances, regarded in Lavoisier's time as elementary, 
have since been shown to be compounds. Thus, quicklime was a 
simple substance until Davy, in 1808, prepared the metal calcium 
and showed that quicklime was the oxide of this metal. Hence, 
we do not say that the substances regarded as simple cannot be 
decomposed, but only that they are substances which we "are not 
able" (at present) to decompose. 

The phrase "at will" is also important. Radium (q.v*) cannot 
be decomposed at will, but it undergoes continuous "disintegra- 
tion," producing the elements helium and lead. We can neither 
hasten, retard, nor stop this spontaneous decomposition. 

The highly interesting experiments of Collie, Paterson, and 
Masson (Chemical Society, London, Annual Report, 1914, pp. 41- 
47) seem to show that the elements helium and neon can be pro- 

* Contraction for quod vide, which see. This abbreviation is used when 
subjects not yet discussed are mentioned. For such subjects, consult the 
index. 



16 COLLEGE CHEMISTRY 






duced by electrical discharges in vacuum tubes and even in a 
closed tube surrounding the vacuum tube. Before long, there- 
fore, the decomposition of elementary substances, and the forma- 
tion of some elements from others, at will, may be a recognized 
possibility, and the foregoing definition may have to be radically 
revised. 

Elements. — The word element is used in two senses. It is 
applied to the simple substance. Thus we speak of "the element 
iron," meaning the metal iron. It is applied also to the iron- 
matter contained in ferrous sulphide or in ferric oxide. The 
reader should note that it is: correct usage to speak of the element 
iron and the element sulphur in ferrous sulphide, but a chemist 
would never say that this compound contained the simple sub- 
stances iron and sulphur. If he did, we should understand him to 
mean that it was a mixture, and we should expect parts of the 
material to be magnetic like iron, and other parts to be yellow and 
soluble in carbon disulphide, which is not the case. In the same 
way the name of an element (such as iron) is applied both to 
the material in combination and to the free substance. Thus 
"iron" may mean free, uncombined, metallic iron, or iron-matter 
in some compound. The sense in which the word is employed 
must be inferred from the context or circumstances. When a 
chemist speaks, as he sometimes does, colloquially, of "iron" in a 
drinking water, for example, we know at once that he refers to 
iron in the form of some compound, for metallic iron does not 
dissolve in water and, if it did, would quickly turn into rust or 
some other form of combination. 

The word element, then, means one of the simple forms of 
matter, either free or in combination. 

In formally describing a body or specimen, the chemist always 
avoids the ambiguity just referred to by naming the components, 
i.e., the substance or substances it contains. He assumes that the 
nature and constituents of these substances will be known to any- 
one hearing or reading the description. If he says the body con- 
tains zinc and sulphur, it is understood that the body is a mixture 
of these simple substances. If it contained these elements in 
combination, the chemist would report that it was sulphide of 
zinc. 



CHEMICAL CHANGE AND METHODS OF STUDYING IT 17 

The Common Elements, — Thousands of different com- 
pound substances are known but, when they are decomposed, it is 
found that the number of different elements contained in them is 
not great. Dozens of substances contain iron, hundreds contain 
sulphur, thousands contain oxygen. In fact, by combining a 
limited number of simple substances, two, three, or four, together, 
in varying proportions by weight, an almost unlimited number of 
different compound substances could be produced. 

A list of the elements appears on the inside of the cover, at the 
end of this book, and contains about eighty names. Of these, a 
large number are rare, and seldom encountered. More than 99 
per cent of terrestrial material is maue up of eighteen or twenty 
elements and their compounds. Only about twenty elements 
occur in nature in their simple, uncombined condition. Three- 
fourths of the whole number are found in combination exclusively, 
and must be liberated by some chemical action. 

Taking the atmosphere, all terrestrial waters, and the earth's 
crust, so far as it has been examined, F. W. Clarke has estimated 
the plentifulness of the various elements. The first twelve, with 
the quantity of each contained in one hundred parts of terrestrial 
matter, and constituting together 99 per cent, are as follows: 

Oxygen. . . 49.85 Calcium . . . .3.18 Hydrogen . . 0.97 

Silicon . . . 26.03 Sodium 2.33 Titanium . . 0.41 

Aluminium . 7.28 Potassium. . . . 2.33 Chlorine . . . 0.20 

Iron .... 4.12 Magnesium . . .2.11 Carbon . . .0.19 

Thus oxygen accounts for nearly one-half of the whole mass. 
Silicon, the oxide of which when pure is quartz and in less pure 
form constitutes ordinary sand, makes up half of the remainder. 
Valuable and useful elements, like gold, silver, sulphur, and mer- 
cury, are among the less plentiful which, all taken together, furnish 
the remaining one per cent. 

Law of Definite Proportions. — In the decomposition of 
mercuric oxide (p. 14) we find that, for every 100 parts of mercury 
liberated, almost 8 parts of oxygen (more exactly, 7.97 parts) by 
weight are set free. Using the numbers commonly employed in 
chemistry, which represent the same proportion by weight: 

Mercuric oxide (216.6) -» Mercury (200.6) + Oxygen (16). 



18 COLLEGE CHEMISTRY 

We find also that mercury and oxygen can be made to combine to 
form mercuric oxide, and the proportions by weight required are 
the same. Moreover, every sample of mercuric oxide, whether 
made by combination, or in any of the other possible ways, always 
contains this proportion of the two elements. We have already 
seen that the oxides of lead and tin contain fixed proportions (p. 9) 
of the metal and oxygen and that ferrous sulphide has a constant 
composition by weight. The same principle is found to apply to 
all chemical compounds, and is stated in the law of definite or 
constant proportions: In every sample of any compound substance, 
formed or decomposed, the proportion by weight of the constituent 
elements is always the same. (For the only known exception to 
this law, see radium.) 

Conservation of Mass. — The most painstaking chemical 
work seems to show that, if all the substances concerned in a 
chemical change are weighed before and after the change, there is 
no evidence of any alteration in the quantity of matter. The two 
weights, representing the sums of the constituents and of the prod- 
ucts, respectively, are, indeed, never absolutely identical, but 
the more careful the work and the more delicate the instrument 
used in weighing, the more nearly do the values approach identity. 
We are able to state, therefore, that the mass of a system is not 
affected by any chemical change within the system. 

This statement simply means that the great law of the conserva- 
tion of mass holds true in chemistry as it does in physics. Chemi- 
cal changes, thoroughgoing as they are in respect to all other 
qualities, do not affect the mass; an element carries with it its 
weight, entirely unchanged, through the most complicated chemi- 
cal transformations. 

Superficial observation, as of a growing tree, might seem to give 
evidence of the very opposite of conservation of matter. But 
here the carbon dioxide gas in the air, the most important source 
of nourishment for plants, is overlooked. Similarly, the gradual 
disappearance of a candle by combustion seems to illustrate the 
destruction of matter. But if we catch the gases which rise through 
the flame (Fig. 12), we find that the gases weigh even more than 
the part of the candle which has been sacrificed in making them. 
When we take account of the weight of the oxygen obtained from 



CHEMICAL CHANGE AND METHODS OF STUDYING IT 19 

the air which sustains the combustion, we find that there is really 
neither loss nor gain in weight. If we carry out chemical changes 
in closed vessels (Fig. 13), which permit neither escape nor access 
of material, we find that the weight does not alter. 





Fig. 12. 



Fig. 13. 



Specific Physical Properties. — It will be seen that, to the 
chemist, knowing the physical properties of all substances is very 
important. By means of the properties, he recognizes and de- 
scribes all the bodies he studies. It may be well, therefore, here 
to give a list of the more important properties, most of which have 
been mentioned in connection with the illustrations we have used. 

In the case of solids, the chief physical properties the chemist 
uses are color, crystalline form, solubility or non-solubility in water 
and occasionally other liquids, the temperature at which the sub- 
stance melts (melting-point), and the density. 

In the case of liquids, he notes the temperature at which the 
liquid boils (boiling-point), the density, the mobility, the odor, and 
the color. 

Finally, in the case of gases, the properties commonly mentioned 
are the color, taste, and odor, the density, solubility in water, and 
the ease or difficulty with which the gas can be liquefied. 



Attributes and Conditions. — There are other qualities 
which a body may possess that we are liable to confuse with the 
specific properties. Thus, the weight of a piece of sulphur is not 



20 COLLEGE CHEMISTRY 

a property of sulphur. A hundred pieces of as many different 
substances might all have the same weight, so that a particular 
weight (say 2 grams) is not a property of any one species of matter. 
Weight, dimensions, and volume are attributes of a body. They 
have different values for different bodies, even when those bodies 
are all composed of the same substance. The attributes are 
physical in nature. They are of great importance in chemistry, 
however, because they afford the only means we have of measuring 
quantities of substances. 

There are still other qualities which a body (or specimen of 
matter) may possess. It has, for example, a certain temperature, 
pressure (state of compression), motion, or electric charge, and it 
may be in solution in some liquid. A body may change in tem- 
perature, pressure, or state of electrification, or it may be dissolved 
in water, or be recovered by evaporation of the liquid, and yet be 
the same specimen. A hundred specimens of as many different 
substances may all have the same temperature — this is not a 
specific property. These are all spoken of as conditions. They 
are physical conditions. In chemistry, conditions are often altered 
in order to bring about, or to stop chemical change, or to modify 
the speed with which it takes place. Thus, we heated the lead 
(raised its temperature) in order to hasten the process of rusting. 
If a substance, or mixture, is capable of undergoing ordinary 
chemical change, then the change is always hastened by raising 
the temperature, and is always delayed or prevented by lowering 
the temperature. Similarly, changing the pressure in a gas, or 
dissolving a substance in some liquid, frequently hastens or de- 
lays a chemical change in which the substance takes part. The 
proper physical conditions are, therefore, considered in connection 
with every chemical operation. Conditions are used to modify 
chemical change. 

Physics in Chemistry. — It will be seen that one cannot 
accomplish anything in chemistry without acquiring and using 
some knowledge of physics. We measure quantities by means of 
the physical attributes, weight and volume. We produce chemi- 
cal change by arranging the physical conditions, for example, by 
mixing, heating, or using an electric current. Physical means are 
the only means we possess for producing, stopping, or modifying 



CHEMICAL CHANGE AND METHODS OF STUDYING IT 21 

chemical changes. Again, we ascertain whether a chemical change 
has taken place or not by observing the physical properties of 
the materials before and after the experiment. Thus, we noted 
that the red, powdery oxide of mercury, when heated, gave a 
liquid metal and a gas. All the phenomena of chemistry are 
physical. A phenomenon is literally something that is seen or, 
more generally, something that affects any of the senses. Observ- 
ing physical phenomena is, therefore, our sole means of studying 
chemical changes. Chemical work is, in fact, entirely dependent 
upon the skilful use of physical agencies, and upon the close obser* 
vation of physical phenomena for its success. 

It is only the inference, following the experiment and the obser- 
vation, that is strictly chemical. If one substance gives two 
different substances, or if two substances give one different sub- 
stance, for example, we infer that a chemical change has occurred. 
We then try to recognize the substances by their properties and 
name them. 

Changes like that of ice into water, or of water into steam, and 
vice versa, are not regarded as chemical changes. These are called 
changes of state, or of state of aggregation. The solid, liquid, and 
gaseous forms are different states of the same substance. 

Law: Explanation: Scientific Method. — There is a 
widely spread impression that a science, like chemistry, is a part of 
the natural order of the universe. It is thought that we are try- 
ing to find the boundaries of chemistry, as they have been pre- 
determined by nature, and to discover the facts, relations of 
facts, and laws which nature has provided as a means of classifying 
the content of the science. Now, the situation is precisely the 
reverse of this. Nature provides only the materials and the phe- 
nomena, and man is attempting to classify them. He divides the 
whole into groups, such as physics, chemistry, botany, etc. Then 
he classifies the facts within each group, in order that he may 
more easily remember them and perceive their relations. He 
often finds that, when new facts are discovered, parts of the 
classification have to be changed. Thus, as we have just seen, 
changes of state are usually assigned to physics, but Ostwald at 
one time suggested that they should be considered as chemical 
phenomena. 



22 COLLEGE CHEMISTRY 

In the preceding pages, we have discussed some of the ways that 
have been invented for classifying the materials and facts assigned 
to chemistry. Thus, we pick out a number of facts of a like 
nature and try to make a single statement which will cover all 
these facts. For example, we find about one hundred thousand 
different substances and, in the case of each substance, every speci- 
men that we have examined contains the same proportions of 
the constituent elements. So we formulate the law of constant 
proportions. A law or generalization in chemistry is a brief state- 
ment describing some general fact or constant mode of behavior. 
We must remember, however, that laws are only true so long as no 
facts in conflict with them are known. There are no laws in nature. 
Nature presents materials and phenomena as she pleases. The 
laws are parts of science, which is made by man, and is a description 
of natural facts as man knows them. As we have seen (p. 18), 
at least one undoubted exception to the law of constant propor- 
tions has recently (1914) been discovered, and other exceptions to 
this law will undoubtedly present themselves. 

One section (p. 5) was entitled: " Explanation of rusting." 
If that paragraph be now re-read, it will be found that, in the 
ordinary (as distinct from the scientific) sense of the word, no 
explanation was given! When we ask a man to "explain" some 
feature in his conduct, we recognize that he might have chosen to 
act otherwise, and we wish to know why he acted precisely as he 
did. Nature, however, has no free will, and cannot tell why she 
presents certain phenomena, and not others. 

On examining the explanation, we find that it simply shows 
that when iron rusts it combines with oxygen from the air. This 
is an additional fact. It shows how iron rusts, namely by taking 
up oxygen, but not why it is able to unite with oxygen. We 
simply do not know why iron can combine with oxygen gas and 
platinum cannot. 

Explanations in chemistry are of three kinds. (1) We usually 
try to show that the phenomenon is not an isolated one. Thus, 
we show that other metals rust. This reconciles us to some ex- 
tent to the fact that iron rusts, and we feel some mental satisfac- 
tion. This is the method of showing that the fact to be explained 
is a member of a large class of similar facts. (2) Next, we try to get 
more information about the fact to be explained. Thus, when, to the 



CHEMICAL CHANGE AND METHODS OF STUDYING IT 23 

acquaintance with the outward manifestations of rusting, we add 
the further information that there is an increase in weight, and 
that this is due to union of oxygen from the air with the iron, we 
feel increased satisfaction, and say that the fact has been "ex- 
plained." (3) If we are still dissatisfied, and can discover no 
further useful facts, we imagine a state of affairs which, if true, 
would classify the fact or add to what we know about it. This step 
we call explaining by means of an hypothesis. We then devote 
our attention to trying to verify the hypothesis. 

The formulation of laws and the making of attempts to explain 
facts are part of what is called the scientific method. The purpose 
of this method is to convert the subject matter into a science, that 
is, into an organized body of knowledge. 

Summary. — In this chapter we have learned: (1) that, 
while there are many substances, there is a limited number of 
entirely different kinds of matter (elements) ; (2) that, in addition 
to constant physical properties, each substance has a constant 
composition by weight. We have also learned that physical 
properties are utilized in manipulations, like nitration and evap- 
oration, as well as for identifying substances, and that physical 
attributes are used for measuring quantities in chemistry and 
physical conditions for guiding chemical change. Finally, we 
have seen that a science is not a natural, but a manufactured 
product, and that the science of chemistry is still in the 
making. 

Exercises,* — 1. What physical properties are used (a) in 
nitration, (6) in evaporation, (c) in the separation and identifica- 
tion of the products from heating mercuric oxide (p. 14)? 

2. Describe: (a) a red-hot rod of iron, 10 cm. long by 1 cm. 
diameter, weighing 58.5 g.; (6) a solution of 5 g. of sulphur in 
20 c.c. (59 g.) of carbon disulphide at 18° C. In doing so, divide 
the description into attributes, conditions, and properties. 

* The exercises should in all cases be studied with minute care. They not 
only serve as tests to show that the chapter has been understood, but very 
frequently (as in No. 4) also call attention to ideas which might not be ac- 
quired from the text alone, or (as in Nos. 1, 2, 5) assist in elucidating ideas 
given in the text which, without the exercises, might not be fully grasped. 



24 COLLEGE CHEMISTRY 

3. Consider the following materials and state whether, so far 
as you can now judge, each is a single substance or a mixture: (a) 
a candle, (6) a cake of soap, (c) an egg. 

4. What are the two most direct ways of showing a substance 
to be a compound? Illustrate each. 

5. If we say that quicklime contains calcium (p. 15), do we 
mean the element or the simple substance calcium? 

6. What explanation was given, (a) of the disappearance of 
mercuric oxide when heated, (6) of the absence of iron and sulphur, 
as substances, from ferrous sulphide? Which of the three kinds 
of explanation was used in each case? 



CHAPTER III 
OXYGEN 

We cannot do better than begin the more systematic study of 
chemistry with oxygen, for it is a most interesting as well as useful 
substance. It is the active component of the air. We depend 
upon it for life, since in its absence we suffocate, for heat, since 
wood, coal, and gas will not burn without it, and even for light 
where oil, gas, or a candle is used. 

We wish to know with which substances we use in the laboratory 
it can combine, as well as the substances on which it .has no 
action. This information will show us how to work, in future, 
without interference from the oxygen in the air and whether oxygen 
has probably played a part in some experiment or not. 

Let us take up, then, (1) the history of the element, (2) what 
materials contain oxygen (occurrence), (3) how we can obtain it in 
a pure state (preparation), (4) what its specific physical properties 
as a substance are, and (5) what it does, and what it cannot do in 
nature and in the laboratory (chemical properties). The classifi- 
cation of the facts about this, and other substances, under five 
heads is somewhat mechanical, but has the advantage of enabling 
the reader quickly to find any required information. 

History of Oxygen. — The Chinese, in or before the eighth 
century, knew that there were two components in the air, and 
that the active one, yin, combined with some metals, and with 
burning sulphur, and charcoal. They even knew that it could be 
obtained in pure form by heating certain minerals, of which one 
was saltpeter. Leonardo da Vinci (1452-1519) seems to be the 
first European to mention the former fact. Mayow (1669) 
measured the proportion of oxygen in the air and discussed fully 
its uses in combustion, rusting, vinegar-making, and respiration, 
but did not make a pure sample. Hales (1731) made it from salt- 
peter, and measured the amount obtainable, but did not see any 

25 



26 COLLEGE CHEMISTRY 

connection between it and the air! Bayen (Apr., 1774) was the 
first to make it by heating mercuric oxide. Priestley (Aug. 1, 
1774) made it by heating the same substance and quite purpose- 
lessly, as he admits, thrust a lighted candle into it and was de- 
lighted with the extreme brilliance of the flame. He had, however, 
entirely incorrect ideas about its nature, and no notion until a 
year later that it was a component of the air. Scheele, a Swedish 
apothecary, had made it in 1771-2 from no less than seven different 
substances and understood clearly that atmospheric oxygen com- 
bined with metals, phosphorus, hydrogen, linseed oil and many 
other substances. But the publisher did not get his book out 
until 1777, and Priestley is usually credited with the " discovery " 
of the element. Finally, Lavoisier (1777) heated mercury in a 
retort (Fig. 14), the neck of which projected into a jar standing 
in a larger dish of mercury. The air, thus enclosed within the jar 
— ,• and the retort, during twelve days lost 

jf_^X\ f\ one-fifth of its volume. Simultaneously, 
\r"ff"k rec ^ particles of mercuric oxide accumu- 
3jjp3 lated on the surface of the mercury in the 
VT "jPn " retort. The residual gas, nitrogen, no 
|| I longer supported life or combustion. The 

oxide, on being heated more strongly, by 
itself, gave off a gas whose volume exactly 
corresponded with the shrinkage undergone by the enclosed air, 
and this gas possessed in an exaggerated degree the properties 
which the air had lost. The proof that oxygen was a component 
of the atmosphere was therefore complete. Later, Lavoisier, in 
the mistaken belief that the new element was an essential con- 
stituent of all sour substances, named it oxygen (Gk., acid-producer) . 

Occurrence. — As we have seen, nearly 50 per cent of terres- 
trial matter is oxygen. Water contains about 89 per cent, the 
human body over 60 per cent, and common materials like sand- 
stone, limestone, brick, and mortar more than 50 per cent of this 
element. One-fifth by volume (nearly one-fourth by weight) of 
the air is free oxygen. 

Preparation of Oxygen. — 1. The oxygen of commerce is 
now made chiefly from liquefied air (q.v.*). The liquid oxygen 
* See p. 15, footnote. 



OXYGEN 



27 



boils at — 182.4°, but the nitrogen boils at an even lower tempera- 
ture ( — 194°). Since the liquid air has a temperature of about 
— 190°, somewhat above that of boiling nitrogen, the latter 
evaporates much more freely than does the oxygen. After a 
time, when the remaining liquid is almost pure oxygen (96 per 
cent), the gas coming off is compressed by pumps into the steel 
cylinders (Fig. 15) in which it is sold. In medicine, patients 
suffering from pneumonia or suffocation obtain some relief by 
inhaling it in this form. It is also used in feeding 
flames, instead of air, when intense heat is required 
(see acetylene torch and calcium light) . 

2. Unfortunately, it is difficult to liberate oxygen 
from natural substances. Saltpeter (potassium nitrate), 
for example, which is found in many soils and can 
be dissolved out with water, gives off oxygen (p. 25) 
only when raised to a bright red heat by the Bunsen 
flame or blast lamp. But it gives up only one-third of 
the oxygen it contains (101.1 g. give 16 g. of oxygen). 

3. In practice, we are compelled to use manufac- 
tured substances. Amongst the artificial substances 
are mercuric oxide, expensive but historically inter- 
esting (p. 14), potassium chlorate, perhaps the most 
convenient for laboratory use, and sodium peroxide. 
Potassium Chlorate (q.v.) is a white crystalline substance used, 
on account of the oxygen it contains, in large quantities in 
the manufacture of matches and fireworks. When heated in a 
tube similar to that in Fig. 11, it first melts (351°) and then, on 
being more strongly heated, it effervesces and gives off a very large 
volume of oxygen. Examination shows that the whole of the 
oxygen it contains (39 per cent) can be driven out. The white 
material which remains after the heating is identical with the 
mineral sylvite. To the chemist it is known as potassium chloride. 
The change, together with the weights of the materials, is as follows: 

Potassium chlorate (122.56) — » Potassium chloride (74.56) + Oxygen (48) 




Fig. 15. 



Potassium (39.1) 
Chlorine (35.46) 
Oxygen (48) 



Potassium (39.1) 
Chlorine (35.46) 



A peculiarity of this action is that admixture of manganese 
dioxide (the mineral pyrolusite) increases very markedly the 



28 



COLLEGE CHEMISTRY 



speed with which the decorqposition of the potassium chlorate 
takes place. Hence, in its presence, and it is generally mixed 
with the chlorate in laboratory experiments (Fig. 16), a sufficient 
stream of the gas is obtained at a relatively low temperature 



U 






• 







*^WVV . 


^X 






_ — ^ 


\, 




^\ 



fcg^Sa 



Fig. 16. 

(below 200°, see p. 29). Hales (p. 25) was the first to collect a 

gas over water (Fig. 16), in order that it might be kept unmixed 
With air. 

4. Oxygen can be obtained conveniently 
from sodium peroxide and water by means of 
generators (Fig. 17) similar to the acetylene 
generators used on automobiles. When 
the metal sodium is burned in air, sodium 
peroxide is obtained as a powder. This 
powder, after being melted, solidifies in com- 
pact, solid form, and is sold as oxone. The 
oxone is bought in a small, sealed tin can, 
the ends of which are perforated in several 
places just before use. When the valve (B) 
is opened, so that the oxygen escapes, the 
water, which fills the generator almost to 
the top, enters the can (C) by the holes in 
the bottom and interacts with the oxone. 
When the valve is shut, the gas continues 
to be generated until it has driven the 

water down again below the level of the bottom of the can. 

Sodium peroxide (78) + Water (18) -> Sodium hydroxide (80) + Oxygen (16) 

Sodium (46) Hydrogen (2.016) Sodium (46) 

Oxygen (32) Oxygen (16) Oxygen (32) 

Hydrogen (2.016) 



I 



1 



Fig. 17. 



OXYGEN 



29 






This method is convenient because it works at room temperature 
and can be started and stopped at will. The sodium hydroxide 
produced is very soluble in water and remains dissolved. Note 
that the name of this substance indicates the elements which 
compose it. 

Catalytic or Contact Action. — The influence of manganese 
dioxide in causing the potassium chlorate to decompose more easily 
(p. 27) well deserves notice. The effect is very striking if some 
pure potassium chlorate is melted carefully, to avoid superheat- 
ing, in a wide-mouth flask (Fig. 18). The flask is provided with 
a wide exit tube, from which a 
rubber tube may lead to a bottle 
inverted in a trough filled with 
water as in Fig. 16. A little 
manganese dioxide is contained 
in the upper, closed tube. No 
effervescence of the chlorate 
can be seen at its melting-point 
(334°) — only a little air, ex- 
panded by the heating, issues 
from the tube. When, however, 
the closed tube containing the 
manganese dioxide is rotated 
into a vertical position (see FlG - 18 - 

dotted lines) , and the black powder falls into the chlorate, the oxygen 
comes off in torrents, in consequence of the enormous acceleration of 
the decomposition. As a precaution against inj ury from an explosion, 
it is advisable to wrap the flask in a towel before turning the tube. 

It must also be noted that the manganese dioxide is not itself 
permanently altered. If the material left after the action is 
shaken with water, the potassium chloride dissolves, while the 
dioxide does not. Filtration (p. 12) then enables us to recover 
the latter, and to ascertain that it has been changed neither in 
quantity nor in properties. 

The only effect of the dioxide is to hasten the decomposition of 
the chlorate, which would otherwise be too slow at 200° (p. 28), 
or even at 334° (its m.-p.) to be of any practical value. Sub- 
stances which hasten a chemical action without themselves under- 




30 COLLEGE CHEMISTRY 






going any permanent change are called contact agents, catalytic 
agents, or catalysts. The process is called contact action or catal- 
ysis (Gk., decomposition, not a very fortunate choice of words). 
Such substances are frequently used in chemistry. The addition 
of a suitable catalyst is one of the conditions (p. 20) for carrying 
out actions in which a contact agent is necessary. Many sub- 
stances of this class are secreted by animals and plants and play 
an important part in digestion, fermentation, and other physiologi- 
cal changes. Their presence often enables very complex chemi- 
cal actions to proceed rapidly at rather low temperatures. 

The oxone, mentioned above, always contains a trace of cuprous 
oxide which hastens the action on water. 

Specific Properties of Two Kinds, Physical and Chemical. 

— We have learned that every substance has its own set of specific 
properties. In describing a substance, it is convenient to divide the 
properties into two classes. The list of substances with which the 
given substance can enter into chemical combination, for example, 
we place under specific chemical properties. Relations of the sub- 
stance to any of the varieties of chemical change belong to this class. 

On the other hand, we do not consider melting or boiling to be 
chemical changes, so we place the temperatures at which the sub- 
stance melts (m.-p.) and boils (b.-p.), its color, etc. (for list, see 
p. 19), under specific physical properties. 

Properties of either class may be used for recognizing a substance. 

Specific Physical Properties of Oxygen. — Oxygen resembles 
air in having neither color, taste, nor odor. The density of a sub- 
stance is, strictly speaking, the weight of 1 cubic centimeter (1 c.c). 
In the case of a gas, we frequently prefer to give the weight of 1000 
c.c. (1 liter), at 0° and 760 mm. (1 atmosphere) barometric pressure. 
For oxygen this weight is 1.42900 grams (Morley). The corre- 
sponding weight for air is 1.293, so that oxygen is slightly heavier, 
bulk for bulk, than air (in the ratio 1.105 : 1). Oxygen can be 
liquefied by compression, provided its temperature is first reduced 
below —118°, which is its critical temperature.* The gas is 

* Each gas has an individual critical temperature (q.v.) above which no 
pressure, however great, will produce liquefaction. The farther the tempera- 
ture of a specimen of the gas is below the critical point, the less will be the 
pressure required to liquefy it. 



OXYGEN 



31 



slightly soluble in water, the solubility at 0° being 4 volumes of gas 
in 100 volumes of water (at 20°, 3 : 100). 

The solubility of oxygen in water, although slight, is in some 
respects its most important physical property. Fish obtain oxy- 
gen for their blood from that dissolved in the water. With air- 
breathing animals (like man), the oxygen could not be taken into 
the system, if it did not first dissolve in the moisture contained in 
the walls of the air sacs of the lungs, and then pass inwards in a 
dissolved state to the blood. 

Liquid oxygen, first prepared by Wroblevski, has a pale-blue 
color. At one atmosphere pressure, that is, in an open vessel, it 
boils at — 182.5°. Its density (weight of 1 c.c.) is 1.13, so that it 
is slightly denser than water. By cooling with a jet of liquid 
hydrogen, Dewar froze the liquid to a snow- 
like, pale-blue solid. A tube of liquid oxygen 
is noticeably attracted by a magnet. 

Six Specific Physical Properties of Each 
Gas. — Although every substance has many 
physical properties, we shall mention only 
those which are used in chemical work, with 
occasionally the addition of any peculiar or 
unexpected quality. It will aid the memory 
to recall the physical properties of a gas, if 
we note that, as a rule, only six such proper- 
ties are mentioned: (1) color, (2) taste, (3) 
odor, (4) density, (5) liquefiability, denned by 
the critical temperature, (6) solubility, usually 
in water only. 

Specific Chemical Properties of Oxy- — 
gen. — The chemical properties of pure 
oxygen are like those of atmospheric air, 
only more pronounced. 

Non-metallic Elements. Sulphur, when raised in advance to the 
temperature necessary to start the action, unites vigorously with 
oxygen (Fig. 19), giving out much heat and producing a familiar 
gas having a pungent odor (sulphur dioxide). This odor is fre- 
quently spoken of as the " smell of sulphur," but in reality sulphur 




32 COLLEGE CHEMISTRY 

itself has no odor, and neither has oxygen. The odor is a property 
of the compound of the two. The mode of experimentation can 
be changed and the oxygen led into sulphur vapor through a tube. 
The oxygen then appears to burn with a bright flame, giving the 
same product as before. 

Phosphorus, when set on fire, blazes in oxygen very vigorously, 
forming a white, powdery, solid oxide — phosphorus pentoxide. 
Burning carbon, in the form of charcoal or hard coal, glows bril- 
liantly and is soon burnt up. It leaves an invisible, odorless gas — 
Garbon dioxide. At high temperatures, oxygen combines readily 
with one or two other non-metals (e.g., silicon, boron, and arsenic), 
and to a small extent (1 per cent at 1900°) with nitrogen. It will 
not combine directly with chlorine, bromine, or iodine, although 
oxides of the first and last can be prepared by using other varieties 
of chemical change. With the six members of the helium family 
(q.v.), of which no compounds are known, and with fluorine, oxygen 
forms no compounds. 

Sulphur (32.06) + Oxygen (32) ->Sulphur dioxide (64.06). 
Phosphorus (62.08)+ Oxygen (80) -^Phosphorus pentoxide (142.08). 
Carbon (12) + Oxygen (32) ->Carbon dioxide (44). 

Metallic Elements. Iron, as we have seen, rusts exceedingly 
slowly in air and, even when red-hot, gives hammer-scale, the black 
solid which is broken off on the anvil, rather deliberately. In pure 
oxygen, a bundle of picture-wire, if once ignited, will burn with 
surprising brilliancy, throwing off sparkling globules of the oxide, 
melted bj^ the heat. This oxide is a black, brittle substance, 
identical with hammer-scale, and different from rust (ferric oxide). 
It contains, in fact, a smaller proportion of oxygen than does the 
latter, and is called magnetic oxide of iron. 

Iron (167.52) + Oxygen (64) -4 Magnetic oxide of iron (231.52). 

All the familiar metals, excepting gold, silver, and platinum, 
when heated, combine with oxygen, some more vigorously, others 
less vigorously than does iron. Oxides of the three metals just 
named can also be made, but only by varieties of chemical change 
other than direct combination. 

Compound substances, if they are composed largely or entirely 
of elements which combine with oxygen, are able themselves to 



OXYGEN 33 

interact with oxygen. Usually, they produce a mixture of the 
same oxides which each element, separately, would give. Hence, 
wood, which is composed of carbon and hydrogen with some 
oxygen, when burnt in oxygen, produces carbon dioxide and water 
(oxide of hydrogen) in the form of vapor. Again, carbon disul- 
phide burns readily, giving carbon dioxide and sulphur dioxide, 
just as do carbon and sulphur, separately. Ferrous sulphide gives, 
similarly, sulphur dioxide and magnetic oxide of iron. 

Tests. A Test for Oxygen. — A test is a property which, 

because it is easily recognized (a strong color, for example), or for 
some other sufficient reason, is commonly employed in recognizing 
a substance. 

Oxygen, as we have seen (p. 14), when pure, is recognized by 
the fact that a splinter of wood, glowing at one end, bursts into 
flame when introduced into the gas. Only one other gas (see 
nitrous oxide) behaves similarly. 

The Measurement of Combining Proportions. — In a 

number of condensed statements we have given the proportions 
by weight of the materials combining. It is now desirable that 
we should know how the necessary measurements are made. The 
most exact measurement of the proportions in which the elements 
combine to form compounds involves manipulations too elaborate 
to be gone into here. One or two brief statements, diagrammatic 
rather than accurate, will show the principles, however. 

If we take a weighed — — — 7n] _^ 

quantity of iron in a test- c= ^- 

tube and heat it with more 

than enough sulphur (an 

excess of sulphur), we get 

free sulphur along with the 

ferrous sulphide (pp. 13-14), 

and no free iron survives. 

We may remove the free 

sulphur by washing the solid FlG - 20 - 

with carbon disulphide. The difference between the weights of the 

ferrous sulphide and the iron gives the amount of sulphur combined 

with the known quantity of the latter. 



34 COLLEGE CHEMISTRY 

As an example of the study of the combination of a metal with 
oxygen, we may weigh a small amount of copper in the form of 
powder in a porcelain boat and pass oxygen over the heated metal 
(Fig. 20). If we limit the oxygen, part of the copper may remain 
unaltered; if we use it freely, the excess will pass on unchanged. 
The original weight of the copper, and the increase in weight, 
representing oxygen, give us the data for determining the compo- 
sition of cupric oxide. The data furnished by one rough lecture- 
experiment, for example, were as follows: 

Weight of boat + copper 4 . 278 g. 

Weight of boat empty 3 . 428 g . 

Difference = weight of copper . . .' 0.860 g. 

Weight after addition of oxygen 4 . 488 g. 

Weight without oxygen 4.278 g . 

Difference = weight of oxygen 0.210 g. 

The proportion of copper to oxygen, so far as this one measure- 
ment goes, is therefore 85 : 21. 

The results of quantitative experiments are often recorded in the 
form of parts in one hundred. To find the percentage of each con- 
stituent, we observe that the proportion of copper is 85 : 85 + 21, 
or T W of the whole. That of the oxygen is T Vg- of the whole. Thus 
the percentages are: 

Copper, 106 : 85 : : 100 : x. x = 80.2. 

Oxygen, 106 : 21 : : 100 : x\ x' = 19.8. 

Naturally, the mean of the results of a number of more carefully 
managed experiments will be nearer the true proportion. The per- 
centages at present accepted as most accurate are 79.9 and 20.1. 

In the case of mercuric oxide, we may decompose a known weight 
of the oxide (p. 14), collect the mercury and weigh it, and ascer- 
tain the oxygen by difference. 

The names of the constituent elements in a compound, together 
with the proportion by weight in which they are present, are called 
the composition of the substance. Thus, the composition of cupric 
oxide is copper : oxygen : : 79.9 : 20.1. This is the percentage com- 
position, but other numbers expressing the same proportion (such 
as 63.57 : 16) will serve the purpose. 

All experiments involving measurement, such as those used in 
determining composition, are called quantitative experiments. 



^ 



OXYGEN 35 

Another Quantitative Experiment. — The following will 
show how the combining proportions may be measured when the 
product is a gas, the weight of which must be ascertained. Sul- 
phur burns in oxygen to form sulphur dioxide. A known weight 
of sulphur is placed in a porcelain boat (Fig. 21), which has already 
been weighed. The U-shaped tube to the right contains a solu- 
tion of potassium hydroxide, which 
is capable of absorbing the resulting — o 
gas. The oxygen enters from the left. 
When the sulphur is heated, it burns 
in the oxygen, and the loss in weight 
which the boat undergoes shows the Fia 21 * 

amount of sulphur consumed. The gain in weight of the U-tube 
shows the weight of the compound produced. By subtracting, 
we get the quantity of oxygen. 

In one experiment, the loss in weight of the boat and its con- 
tents (= sulphur) was 1.21 g. The weight gained by the U-tube 
was 2.42 g. The difference (= oxygen) is 1.21. The proportion 
of sulphur to oxygen in sulphur dioxide is therefore 1.21 : 1.21 or 
1 : 1 or, in percentages, 50 : 50. This proportion is very close to 
the accepted value (p. 32), 32.06 : 32. 

The same method could be used for carbon, for the carbon 
dioxide produced would be absorbed in the solution of potassium 
hydroxide. 

Combustion. — Violent union with oxygen is called, in 
popular language, combustion or burning. Yet, since oxygen is 
only one of many gaseous substances known to the chemist, and 
similar vigorous interactions with these gases are common, the 
term has no scientific significance. Even the union of iron and sul- 
phur gives out light and heat, and is quite similar in the chemical 
point of view to combustion. 

A misleading term often used in this connection is kindling 
temperature. It gives the impression that there is a definite tem- 
perature at which combustion will start. But the temperature 
is only one of the conditions which produce combustion. Finely 
powdered iron will start burning at a lower temperature than will 
an iron wire, because it presents relatively more surface to the gas. 
Again, if the oxygen is at less than one atmosphere pressure, the 



36 COLLEGE CHEMISTRY 

wire will require to reach a higher temperature before combustion 
will begin. Finally, the vapor of methyl alcohol and air requires 
to be raised above a red heat before combustion starts, but a pocket 
cigar-lighter sets fire to this very mixture by means of a contact 
agent (a thin platinum wire) without any other means of heating 
being required. So that, the conditions under which combustion 
begins involve the physical condition of the solid, the pressure of 
the gas or vapor, the presence or absence of a contact agent and 
the nature of the contact agent, as well as the temperature. No 
definite kindling temperature can be given, unless the other con- 
ditions are specified also. Kindling conditions involve several 
variables, of which the temperature is only one. 

Oxidation. — The slower union with oxygen which occurs in 
rusting is called oxidation. We shall see later, however, [that it 
has been found convenient to stretch this term so as to cover com- 
binations of other elements than oxygen, and even to include 
actions not involving combination. At this point we can discuss 
only oxidation by oxygen. 

This process of slow oxidation by oxygen, although less con- 
spicuous than combustion, is really of greater interest. Thus the 
decay of wood is simply a process of oxidation whereby the same 
products are formed as by the more rapid ordinary combustion. 
Sewage is mixed with large volumes of river water, the object being, 
not simply to dilute the sewage, but to mix it with water containing 
oxygen in solution. This has an oxidizing power like that of oxy- 
gen gas and, through the agency of bacteria, quickly renders dis- 
solved organic matters innocuous by converting them for the 
most part into carbon dioxide and water. Thus, a few miles 
further down the stream, the water becomes as suitable for drink- 
ing as it was before the sewage entered. In our own bodies we 
have likewise a familiar illustration of slow oxidation. Avoiding 
details, it is sufficient to say that the oxygen, from the air taken 
into the lungs, combines with the haemoglobin in the red blood- 
corpuscles. In this form of loose combination, it is carried by the 
blood throughout our tissues and there oxidizes the foodstuffs 
which have been absorbed during digestion. The material prod- 
ucts are carbon dioxide and water, of which the former is carried 
back to the lungs by the blood, and finally reaches the air during 



OXYGEN 37 

exhalation. The important product, however, is the heat, given 
out by the oxidation, which keeps the body warm. 

The opposite of oxidation, the removal of oxygen, is spoken of 
in chemistry as reduction. But this term, also, has been stretched 
to cover other kinds of chemical change. 

Spontaneous Combustion. — Sometimes a mere slow oxi- 
dation develops into a combustion, which is then known as spon- 
taneous combustion. To understand this, we must note the fact 
that a given weight of material, sslj, iron, in combining with oxy- 
gen to form a given oxide, will liberate the same total amount of 
heat whether the union proceeds rapidly or slowly. If the action 
proceeds slowly, and the material being oxidized is freely exposed 
to the air, the latter will become heated and will carry off the heat 
as fast as it is produced. Thus, no particular rise in temperature 
will occur. If, however, the material is a poor conductor of heat, 
like hay or rags, and there is sufficient air for oxidation, but not 
enough to carry off the heated air, the heat may accumulate and a 
temperature sufficient to start combustion may be reached. Such 
a situation sometimes arises in hay-stacks. It occurs also when 
rags, saturated with oils used in making paints (linseed oil and 
turpentine) are left in a heap. These oils, in "drying," combine 
with oxygen from the air and turn into a tough, resinous material. 
The rags, being poor conductors of heat, finally become hot 
enough to burst into flame, and serious conflagrations often owe 
their origin to causes such as this. Oily rags should always be 
disposed of by burning, or should at least be placed in a closed can 
of metal. Fires in coal bunkers of ships arise from the same cause 
— slow oxidation, with accumulation of the resulting heat. That 
coal does undergo slow oxidation, especially when freshly mined, 
is shown by the fact that such coal, if left exposed to the air for 
months, may lose 10 per cent or more of its heating power. 

Uses of Oxygen. — A number of the practical applications 
of oxygen have already been mentioned. For example, in the 
foregoing section we have referred to its use in breathing, its r61e 
in decay, which is a beneficent process because it removes much 
useless matter which might otherwise cause disease, and its value 
in the disposal of sewage. Power and heat for commercial pur- 



38 COLLEGE CHEMISTRY 

poses are almost all obtained by the burning of coal, in which oxy- 
gen from the air plays a large part. If we had to purchase the 
oxygen as well as the coal, we should require at least three tons of 
oxygen for every ton of coal. 

Oxygen in cylinders and oxygen generators are used to restore 
the supply in the atmosphere of submarine boats, as well as for the 
purposes already mentioned (p. 27). 

Substances Indifferent to Oxygen, — Finally, since the 
atmosphere contains so large a proportion of oxygen, substances 
which do not oxidize and, when heated, do not burn, have many 
uses. Gold, silver, and platinum are of this kind (p. 32), and are 
used for ornaments. The last is used for crucibles in which bodies 
are heated in the laboratory. Although iron burns in pure oxygen, 
it does not oxidize rapidly in the air even when heated, and so is 
used for making vessels for cooking and in constructing fireproof 
buildings. 

Compounds, already fully oxidized, are naturally not com- 
bustible. Of this nature are sandstone, granite, brick, porcelain, 
glass, and water. All these are, therefore, fireproof. Moreover, 
these substances do not give off oxygen when heated (water de- 
composes slightly). Glass and porcelain thus neither lose nor 
gain in weight when heated, and are suitable materials for labora- 
tory apparatus. 

Activity and Stability. — A substance which enters into 
combination vigorously, as does oxygen, is called chemically 
active. Nitrogen, on the other hand, is relatively inactive. An 
active element, since it combines eagerly, naturally holds tena- 
ciously to the matter with which it has combined. An active ele- 
ment implies, therefore, also one which is in general difficult to 
liberate from combination. Its compounds are in general rel- 
atively stable. Thus, many oxides, and the natural compounds 
just mentioned (sandstone, granite, brick and porcelain, the last 
two made from clay), do not lose oxygen even at a white heat and 
are very stable. 

Exercises. — 1. What percentage by weight of free oxygen is 
obtained by heating: (a) mercuric oxide, (b) potassium nitrate, 



OXYGEN 39 

(c) potassium chlorate? At $1.50 (7/8), $0.15 (8d), and $0.15 (8d) 
per kilogram, respectively, which is the cheapest source of oxygen? 

2. Using the data on pp. 30-31, calculate the weight of oxygen 
dissolved by 1000 c.c. of water at 0°. 

3. Why does a forced draft make a fire burn more rapidly? 

4. Why does a naked flame sometimes cause an explosion in a 
mine, when the air of the mine is filled with coal dust? 

5. The substances, like phosphorus and sulphur, which burn 
rapidly in ordinary oxygen, combine very, very slowly with oxygen 
which has been freed from moisture by careful drying. How is 
this effect of water to be classified? 

6. Air is 20 per cent oxygen. Why does iron burn brilliantly 
in pure oxygen, but not in air? 



CHAPTER IV 

ATOMIC WEIGHTS, SYMBOLS, FORMULA, AND 
EQUATIONS 

We have repeatedly called attention to the quantities of the 
substances taking part in chemical changes, and particularly to 
the constant relation between the weights of each element in a 
given substance (pp. 17-18) . This matter is of great importance in 
chemistry. If a cargo of copper ore is to be purchased, we do not 
wish to pay for the rock that all specimens of the ore contain in 
larger or smaller proportion. So we secure a fair sample of the ore 
and have an analysis made by a chemist. The analysis, in this 
case, is a measurement of the proportion of the valuable metal in 
the sample. The price will then depend largely upon the propor- 
tion of the copper per ton of ore. The making of analyses — that 
is, chemical measurements — plays a very large part in all in- 
dustries which involve the consumption or manufacture of ma- 
terials. Quantitative measurements, aside from their theoretical 
interest, are therefore of the greatest practical importance. Hence 
we must now discuss them once more. 

The Compositions of Substances. — Our present purpose 
is to compare the proportions by weight of the elements composing 
several compounds, in order to see whether the numbers are really 
as irregular as, in the examples we have heretofore given, they 
have appeared to be, or whether there is any way of relating and 
simplifying the numbers. 

In order to have a fair sample of these proportions, we shall in- 
clude the compositions of a few substances for which the data 
have not yet been given. Potassium hydroxide (p. 35) has the 
composition: potassium (a metal) 39.1, oxygen 16, hydrogen 
1.008, in a total of 56.108 parts. Water (oxide of hydrogen) con- 
tains: oxygen 16 and hydrogen 2.016 parts by weight. When 
iron burns in chlorine, which is a yellow gas, it gives ferric chloride 

40 



ATOMIC WEIGHTS, SYMBOLS, FORMULA, AND EQUATIONS 41 

with the proportions: iron 55.84, chlorine 106.38. When ferric 
chloride is heated in a stream of hydrogen gas, a part of the 
chlorine is removed, and ferrous chloride remains: iron 55.84, 
chlorine 70.92. 

To make the comparison easy, we have limited the number of 
substances to five of those previously discussed, together with the 
four just mentioned, and have also arranged the proportions in 
the form of a table. 



PROPORTIONS BY WEIGHT OF THE ELEMENTS 

COMPOUNDS 


IN CERTAIN 


Name of Compound. 


Iron. 


Oxy- 
gen. 


Sul- 
phur. 


Potas- 
sium. 


Chlo- 
rine. 


Hydro- 
gen. 


Ferric oxide (p. 9) 

Ferrous sulphide (p. 14) ... . 
Potassium chlorate (p. 27) . . . 

Sulphur dioxide (p. 32) 

Iron oxide (magnetic) (p. 32) . . 

Potassium hydroxide 

Water 


111.68 
55.84 

167 "52 

55\84 
55.84 


48 

'48' 
32 
64 
16 
16 


32^06 
32! 06 


39 
39 


i 

i 


35 

106 
70 


46 

38 
92 


i 

2 


008 
016 


Ferric chloride 

Ferrous chloride 




Atomic weights 


55.84 


16 


32.06 


39.1 


35.46 


1.008 



Study of the Foregoing Table. — When we first examine 
the numbers in the horizontal lines of the table, we observe that 
the numbers, with the exception of those for oxygen, all involve 
decimal fractions. From this we infer that whole numbers must 
have been chosen intentionally for oxygen. This is, in fact, the 
case. When we next look down the oxygen column, we observe 
that 48 = 3 X 16 and 32 = 2 X 16 and 64 = 4 X 16. All the 
oxygen weights are multiples of 16 by some integral (whole) num- 
ber. In the hydrogen column, the same regularity appears, for 
2.016 = 2 X 1.008. Following up this idea, we find in the iron 
column, 55.84 occurring thrice, and discover that 111.68 = 2 X 
55.84, and that 167.52 = 3 X 55.84. Similarly, in the chlorine 
column, the numbers are multiples of 35.46 by unity or some other 
integer. Thus, the proportion of each element, in various com- 
pounds, can be represented by a fundamental number — a sort of 
unit quantity — multiplied when necessary by the proper integer. 



42 COLLEGE CHEMISTRY 

Now this rule is not confined to these nine compounds, involving 
only six different elements. If we provided a column for every 
known element (about eighty would be needed), and entered the 
composition of every known compound, we should find the same 
rule to hold. This rule can be stated as follows: 

Law of Combining Weights. — In every compound sub- 
stance, the proportion by weight of each element may be expressed 
by a fixed number, a different one for each element, or by a mul- 
tiple of this number by some integer (whole number) . 

Since the proportion by weight in which two (or more) elements 
combine is a chemical property, this is a chemical law. Clearly, it 
does not apply to mixtures, for any irregular proportion could be 
used in the physical process of mixing. 

Explanation of this Law, Atoms and Atomic Weights. — 

To explain this law it was necessary to use the third kind of ex- 
planation (p. 23), namely the making of an hypothesis. The 
details of how two substances combine cannot be seen, so chemists 
had to imagine some details which would account for the possession 
of an individual unit weight by each element. If oxygen is com- 
posed of minute, invisible particles, which are all alike in weight, 
and hydrogen and potassium are of the same nature, except that 
the weight of the particle of each kind of element is different, we have 
the basis of an explanation. We have to suppose, further, that, 
when elements combine, the particles adhere in pairs or groups, as 
wholes, and are never broken. In this way the particle of each 
variety of elementary matter will have a definite, unchangeable 
weight, which will be one of its fixed properties. If the relative 
weights of the particles of oxygen, potassium, and hydrogen are in 
the proportion of the combining numbers in the table, namely 
16 : 39.1 : 1.008, the whole situation becomes clear. Chemical 
union must consist, in detail, in the union of the particles of the 
elements to form the particles of the compound. For each particle 
of potassium hydroxide, one particle each of the three elements 
is required. 

For each particle of water, where the proportion of oxygen to 
hydrogen is 16 : 2.016, evidently one particle of oxygen and two 
particles of hydrogen are necessary. Varying, intermediate pro- 



ATOMIC WEIGHTS, SYMBOLS, FORMULAE, AND EQUATIONS 43 

portions are impossible, because the particles of the elements are 
permanent, are never broken, and combine as wholes, and in a 
uniform way through the mass. The only possible variation 
would be to take different relative numbers of the particles — for 
example, two of oxygen to two of hydrogen (2 X 16 : 2 X 1.008). 
But this product would have a different composition from water, 
and would not be water. This compound, with the double pro- 
portion of oxygen, is indeed known (it is hydrogen peroxide), and 
is the only other known compound of these two elements. 

This theory fully explains why the combining proportions of 
each element, in different compounds, can always be expressed 
by a fixed, unit number (which represents the weight of the 
ultimate particle of that element), multiplied, when necessary, 
by a whole number (representing the number of particles of 
the element required to form a particle of the compound in 
question) . 

This explanation was first offered by Dalton, a schoolmaster 
of Manchester in 1802. Borrowing an idea from the speculations 
of the Greek philosophers, he called the particles of elements 
atoms (Gk., not cut, or not divided). The atoms of any one element 
are all alike in weight, as well as in other properties, but the atoms 
of different elements differ in weight. 

The particles made by uniting two or more atoms, as in forming 
a particle of a compound, are called molecules (Gk. ; a little 
mass) . 

A chemical combination of two simple substances consists, then, 
in an elaborate re-grouping of the atoms of both elements so that 
molecules of the compound are formed. Definite proportions by 
weight are required, in order that the atoms of each element may 
be available in the correct proportion, 1 atom : 1 atom, or 1 : 2, 
or 2 : 3, or in some similar, usually simple ratio. 

The result was called the atomic theory. For long it remained 
an hypothesis, or sort of guess. Recently, however, we have 
obtained independent proof that molecules and atoms are real 
(see Radioactivity) , for we can now count and measure the weight 
of individual molecules, and we even know something of the inside 
structure of atoms. 

The fundamental numbers, one for each element, being the 
relative weights of the atoms, are called atomic weights. 



44 COLLEGE CHEMISTRY 

Symbols and Formulae. — One self-evident use for the 
atomic weights is in stating the compositions of compounds. To 
make the statement as simple as possible, symbols, first used by 
Berzelius, represent the atomic weight of each element. Thus, H 
stands for 1.008 parts, or 1 atom, of hydrogen, and for 16 parts, 
or 1 atom, of oxygen. When several elements have the same 
initial letters, another letter is added: C for one atomic weight of 
carbon, Ca for one atomic weight of calcium, CI for 35.46 parts of 
chlorine. When the names of the elements are not alike in all 
languages, the symbol is frequently based on the Latin name, as 
Fe for iron (ferrum) and Pb for lead (plumbum), or on the German, 
as K for potassium (kalium). The symbols are international. A 
list of the elements, with their symbols and atomic weights, is 
printed inside the back cover of this book. 

The composition of any compound can thus be stated by setting 
down the necessary symbols, together with the whole numbers, if 
any, by which the atomic weights are multiplied. The result is a 
formula. For example, ferric oxide contains iron 111.68 and oxy- 
gen 48 parts (p. 41). This is equivalent to iron 2 X 55.84 and 
oxygen 3 X 16. This again is equivalent, in symbols, to 2 X Fe 
and 30. The formula is written Fe 2 3 . Ferrous sulphide is a 
simpler case: iron 55.84 and sulphur 32.06, or, in symbols, FeS. 
The reader should now examine the whole table on p. 41, and 
work out the formula of each compound and write it in the 
margin. 

Equations. — It is now possible to abbreviate the condensed 
statements we have been using to represent the substances and 
their quantities in chemical reactions. Thus, the three statements 
on p. 32, when translated into symbols, are as follows: 

S + 20 -> S0 2 . 
2P + 50-+P 2 5 (P = 31.04). 
C + 20->C0 2 (C = 12). 

When no coefficient appears before or after a symbol, 1 is to be 
understood. 

Much practice is required to enable one to make and under- 
stand equations. The reader should therefore at once turn back 
to the statements on pp. 9, 14, 17, 27, and 28, obtain the necessary 



ATOMIC WEIGHTS, SYMBOLS, FORMULAE, AND EQUATIONS 45 

atomic weights and symbols from the table at the end of the book, 
and construct the equation in each case. 

The term " equation'' refers to the fact that the total weight 
of matter on both sides is always the same. In other respects, 
such as in the nature of the substances, the two sides are entirely 
different. 

Derivation of Formulse from Experimental Data. — In 

the condensed statements referred to (by page) in the foregoing 
section, the numbers given were already multiples of the atomic 
weights, and the formulse were therefore easy to make. It re- 
mains to show how the formula may be constructed from the 
weights obtained in an experiment. 

In the quantitative experiment on the composition of cupric 
oxide (p. 34), the proportion found was: copper 85, oxygen 21. 
In the formula, the same proportion is to be expressed by means 
of multiples of the atomic weights. If we divide each of these 
numbers by the corresponding atomic weight, the quotient will 
be the number by which the atomic weight must be multiplied. 
The atomic weights are Cu = 63.57, = 16. 85 ^-63.57 = 1.3, 
and 21 -f- 16 = 1.3. The proportion of copper to oxygen in the 

,85 . 63.57 X 1.3 

compound, — , now becomes • 

2iY lo X l.o 

But this proportion must be expressed in multiples of the atomic 

weights by whole numbers. Dividing above and below by 1.3, we 

. 63.57 X 1 

get ^6^' 

Now the symbols stand for the atomic weights. Substituting 

the symbols, the proportion becomes -~ . The formula is, 

u x i 

therefore, CuO. 

Applying the same process to the case of sulphur dioxide (p. 35) : 

Sulphur 32.06 32.06 X 1 _ S X 1 
Oxygen "32 16x2 "Ox2'° r 2 * 

If the composition of the substance has been stated in percent- 
ages, the same device is used. Thus, the case of sodium sulphate 
works out as follows: 



46 



COLLEGE CHEMISTRY 



Elements. 


Percentages. 


At. Wt. Quotient -f- Formula. 


Sodium 

Sulphur 

Oxygen 


32.43 
22.55 
45.02 


23 X 1.41 0.705 NaX 2 
32 X 0.705 0.705 S 
16 X 2.814 0.705 0x4 



The formula is, therefore, Na 2 S04. 

It is obvious that, after we have found out what elements com- 
pose a given compound, we are still unable to write its formula. 
We may not simply set the symbols down, side by side. A meas- 
urement must be made, in order that we may find out the factors 
by which the atomic weights are to be multiplied. 



Answers to Some Questions. — Why was a whole number 
assigned to oxygen? Oxygen was chosen as the basis of the system 
because the exact determinations of the combining weights of most 
of the elements have actually been made by direct union with 
oxygen, or with the help of but one intermediate step. If the 
question had been one of mathematics, hydrogen, the element 
with the lowest combining proportions, would have furnished the 
basis and unit of the scale. But the question was the practical 
one of getting the most accurate measurements for the relative 
magnitudes of the numbers. Hydrogen combines with only a few 
of the elements, and the proportion of hydrogen is usually so small 
that the weights of this element cannot be measured so accurately 
as can the much larger weights of oxygen and of the other elements. 
So oxygen was selected as the basal element. 

Why was 16 assigned to oxygen, rather than 1 or 100, or some 
other whole number? The number 16 was chosen in order that 
the advantage of having a mathematical unit, or something close 
to it, in the scale, might be retained also. With this value, hydro- 
gen became 1.008. A whole number smaller than 16 would make 
the atomic weight of hydrogen less than unity. With H = 1, 
the value for oxygen becomes about 15.9, and the values for all the 
elements are changed in proportion. The result of such a change 
would be that the values for the common elements would not be so 
close to whole numbers as they at present are (e.g., C = 12.00, 
N = 14.01, Na (sodium) = 23.00, K = 39.1, P = 31.04). With 



ATOMIC WEIGHTS, SYMBOLS, FORMULAE, AND EQUATIONS 47 

O = 16, it is possible, and of course more convenient, in many 
cases to use the nearest whole number in ordinary calculations. 

The answers to the two foregoing questions show why the scale 
of the numbers was fixed as it is. Of course, multiplying or 
dividing all the atomic weights by any number, whole or fractional, 
would not affect their scientific accuracy. The choice of scale is 
merely a matter of convenience. 

In physics there is one unit of weight, the gram, for all kinds of 
matter. Is it the case that in chemistry a different unit of weight 
is employed for each element? This is the exact situation, and 
it is one peculiar to chemistry. It does not represent an arbitrary 
decision of the chemist, however. It is due to the fact that the 
atoms of any one element have the same weight, but that the 
atoms of different elements have different weights. The atom of 
uranium is 238 times as heavy as that of hydrogen, and its com- 
bining proportions, therefore, are in general greater in the same 
ratio, while the atoms of the other elements have weights falling 
between these limits. 

There is still one question to be asked. Why take 16 for oxy- 
gen rather than 8 or 32? In other words, may we not multiply or 
divide any one (or more) of the individual atomic weights by a 
whole number? The answer is that, thus far, we have not met 
with any reason for not doing so. With = 8, and H still 1.008, 
the composition of water would be represented by the formula HO 
instead of H 2 (where = 16). In a later chapter (Chap. VIII), 
however, we shall see that the individual numbers actually chosen 
meet certain other conditions, in addition to those already men- 
tioned, and are on that account preferable to any other set. 

Law of Multiple Proportions, — We have already met with 
several instances in which two elements combine in more than 
one proportion by weight, and form therefore more than one com- 
pound. Thus two oxides of iron and two chlorides of iron have 
been mentioned (p. 41), and two oxides of hydrogen, water and 
hydrogen peroxide, are known (p. 43). This general fact was dis- 
covered before the law of combining weights (p. 42) had been for- 
mulated, and is a particular case of this law. It was discovered 
by Dalton (1804) and was embodied by him in a statement known 
as the law of multiple proportions, which ran somewhat as follows : 



48 COLLEGE CHEMISTRY 

If two elements unite in more than one proportion, forming two or 
more compounds, the quantities of one of the elements, which in the 
different compounds are united with identical amounts of the other, 
stand to one another in the ratio of integral numbers, which are 
usually small. 

The two chlorides of iron illustrate the law. Ferric chloride 
contains iron 55.84 and chlorine 106.38, and ferrous chloride iron 
55.84 and chlorine 70.92. Thus the quantities of chlorine united 
with identical amounts of iron (namely, 55.84 parts) stand in the 
ratio 106.38 : 70.92, or 3 : 2. 

Exercises. — 1. From the data on p. 9 and the atomic weights, 
calculate the formula of lead oxide. Construct also the equations 
for the decomposition of potassium chlorate (p. 27), and for the 
combination of phosphorus and oxygen (p. 32). 

2. When 1 g. of sodium burns in oxygen, it produces 1.7 g. of 
the oxide. What is the formula of the latter and the equation? 

3. If 26 g. of mercurous oxide are required to give, by heating, 
1 g. of oxygen, what is the formula of the substance? 

4. What are the formulae of the substances possessing the fol- 
lowing percentage compositions? 

I II III 

Magnesium, 25.57 Sodium, 32.43 Potassium, 26.585 

Chlorine, 74.43 Sulphur, 22.55 Chromium, 35.390 

Oxygen, 45.02 Oxygen, 38.025 

5. What are the percentage compositions of substances possess- 
ing the following formulae: Mn 3 4 , KBr, FeS0 4 ? 

6. Compare the formula of mercurous oxide, found in 3, with 
that of mercuric oxide, and show how the compounds illustrate 
the law of multiple proportions (p. 48). 

7. If the atomic weight of potassium were 13.03, and the other 
atomic weights were unchanged, what would be the formulae of 
(a) potassium hydroxide, and (b) potassium chlorate? 



CHAPTER V 

HYDROGEN 

Having learned something of the nature of the atmosphere, 
and particularly of oxygen, its most active component, we turn 
now to water, a substance as closely connected with our daily life 
as is air. We find that it is a compound of oxygen and hydrogen, 
and the latter element, therefore, may be taken up next. Hydro- 
gen is of interest on its own account because it is often used in 
filling balloons, and nearly half the bulk of ordinary illuminating 
gas is free hydrogen. 

History. — That hydrogen is a distinct kind of gas was first 
established by Cavendish (1766). Somewhat later (1781), he 
showed that, when it burned in the air, it gave a vapor which 
could be condensed to liquid water. Since oxygen was then known 
to be the substance with which combustibles united, this proved 
that water was a compound of hydrogen (Gk., water producer) and 
oxygen. 

Occurrence. — Free hydrogen is found, mixed with varying 
proportions of other gases, in exhalations from volcanoes, in 
pockets found in certain layers of the rock-salt deposits, and in 
some meteorites. The air contains not over 1 part in 1,500,000. 
The lines of hydrogen are prominent in the spectra of the sun and 
of most stars. 

In combination, it constitutes about 11 per cent of water. It is 
an essential constituent of all acids. It is contained also, in com- 
bination with carbon, in the components of natural gas, petroleum, 
and all animal and vegetable bodies. 

Preparation by the Action of Metals on Cold Water. — 

To liberate hydrogen from water, it is necessary to use some ele- 
ment with which the oxygen of the water will combine even more 
eagerly than with hydrogen, and to offer this element in exchange 
for the hydrogen. 

49 











v°° 

0*0 

°c°» 






On 




• 


rS 





50 COLLEGE CHEMISTRY 

The more active metals, such as potassium (K), sodium (Na), 
or calcium (Ca), displace hydrogen rapidly from cold water. Po- 
tassium and sodium are lighter than water, and float on the sur- 
face. In the case of the former, so much heat is liberated that the 
hydrogen catches fire, and with neither metal is the experiment 
safe in the hands of a novice. Calcium sinks to the bottom, and 
acts rapidly, but not violently, so that the gas is easily collected 
(Fig. 22). The pieces of these metals, of course, act upon only a 
small part of the water in the vessel. In each case 
the metal displaces one-half only of the hydrogen in 
that part of the water upon which it acts. The 
other products are the hydroxides of potassium, 
sodium, and calcium, respectively. The two former 
dissolve, leaving a clear liquid when the metal is 
all gone, but may be recovered as white solids by 
evaporation. The calcium hydroxide (slaked lime) 
is dissolved only in part, and much of it may be seen 
suspended in the water after the action. 
An alloy of lead with sodium (35 per cent), sold under the name 
of hydrone, affords a convenient substitute for sodium in the fore- 
going actions. 

The Making of Equations. — To make an equation we must 
have the results of quantitative measurements. These furnish 
us with the composition of each substance concerned. The com- 
position, expressed in multiples of the atomic weights, is recorded 
in the formula for the substance. If we are in possession of the 
necessary formulae, we can write the equation. 

For example, the composition of water is: hydrogen 2 X 1.008, 
oxygen 16. In symbols, this is 2H and O, and the formula is, 
therefore, H 2 0. The composition of potassium hydroxide is: 
potassium 39.1, oxygen 16, hydrogen 1.008, and the formula, there- 
fore, KOH. In calcium hydroxide the proportions are: calcium 
40.07, oxygen 2 X 16, hydrogen 2 X 1.008, and the formula Ca(OH) 2 . 

To make the equation, we first write down the formulas of the sub- 
stances used and produced: 

K + H 2 -> KOH + H. 
Na + H 2 -> NaOH + H. 
Skeleton: Ca + H 2 -> Ca(OH) 2 + H. 



HYDROGEN 



51 



Next we must balance this equation, if necessary. That is, we 
must adjust it so that there are equal numbers of atomic weights 
(or atoms) of each element on both sides of the equation. This is 
necessary only in the third equation, and is done because, accord- 
ing to the law of conservation of mass, there must be the same 
quantity of each element after the reaction as there was before it. 
On examining the third equation, we note that there is 20, in the 
(0H) 2 , on the right side and only O on the left. We therefore 
place a 2 in front of the H 2 0, for we cannot get the additional oxy- 
gen excepting by using more water: 
Balanced: Ca + 2H 2 -» Ca(OH) 2 + 2H. 

The number of atomic weights of hydrogen is made equal by using 
2H on the right side. 

The coefficients in front of a formula multiply the whole formula. 
Thus, 2H 2 is equivalent to 2(H 2 0). A subscript coefficient 
following a symbol, however, multiplies that symbol only. Thus 
H 2 is equivalent to (H) 2 0, or (2 X H + 0). 



an Equation, - 

what substances 



-1. Find out, by 
are used and what 



Four Steps in Making 

observation and experiment 
substances are produced. 

2. Find the formula of each substance used or produced. 

3. Set the formulae down as a skeleton equation, placing the 
formulae of the substances used on the left, and of those produced on 
the right. 

4. Adjust, or balance the equation, if necessary. 
The reader must practice the 

making of equations, until he can 
do it quickly. The text contains 
many equations, but more usu- 
ally only the data required for 
making them (the formulae of the 
substances) are given. 

Hydrogen from Metals and 
Water at a High Tempera- 
ture. — With steam at a red 
heat, metals like iron, zinc, and magnesium interact vigorously. The 
steam, generated in a flask, enters at one end of the tube containing 
the metal (Fig. 23) , and the hydrogen passes off at the other. Since, 




52 COLLEGE CHEMISTRY 

at a red heat, all hydroxides, except those of potassium and sodium, 
are decomposed into an oxide of the metal and water, as, for 
example, Mg(OH) 2 — > MgO + H 2 0, the oxides are formed in this 
case: 

Mg + H 2 -> MgO + 2H. 

Iron gives the magnetic oxide Fe 3 4 . 

Making Equations, Again. — The skeleton equation for the 
action of iron on steam is : 

Skeleton: Fe + H 2 -> Fe s 4 + H. 

We are not permitted to alter these formulae themselves, but we 
may put coefficients in front of any of them to make the number 
of atomic weights alike on both sides. A useful rule is to pick out 
the largest formula and reason back from that. Here, this is 
Fe 3 4 . To get Fe 3 , we must start with 3Fe, and to get 4 , we must 
start with 4H 2 : 

Balanced: 3Fe + 4H 2 -* Fe 3 4 + 8H. 

Acids. — In making hydrogen in the laboratory, the acids are 
used almost exclusively. The common acids are hydrochloric acid 
(HC1, Aq), and sulphuric acid (H 2 S0 4 , Aq). The usual forms are 
mixtures containing water, the variable amount of the latter being 
indicated by the symbol Aq.* The former is a solution of a gas, 
hydrogen chloride. The "pure concentrated" hydrochloric acid 
used in laboratories contains nearly as much of the gas (39 per 
cent by weight) as the water can dissolve. The "commercial" 
acid contains impurities and is also less concentrated. The "con- 
centrated" sulphuric acid is an oily liquid containing practically 
no water. The "commercial" sulphuric acid contains 6 to 7 per 
cent of water, besides impurities. Acetic acid (HC0 2 CH 3 , Aq) is a 
solution of a liquid in water, and is the acid found in vinegar. 

All the "dilute" acids contain 70 to 80 per cent of water. The 
water, as a rule, takes no part in the chemical changes in which the 
acids are concerned, and is therefore omitted from the equations. 

* The formula H 2 stands for a fixed proportion of water, namely 18 parts. 
The water in these solutions is not combined, and can be varied in amount, so 
that the formula H 2 may not logically be employed here. 



HYDROGEN 53 

The name "acid" is restricted to one class of substances having 
certain definite characteristics. Hydrogen is the one essential con- 
stituent of all acids. Their aqueous solutions have a sour taste and 
change the color of litmus from blue to red. When free from water 
they do not conduct electricity. When dissolved in water they 
conduct, and are decomposed by the electric current. In aqueous 
solution, also, their hydrogen (or one unit weight of it in the case 
of acetic acid) is displaced by certain metals. 

Radicals. — In describing the chemical behavior of acids, we 
speak of the hydrogen as the positive radical, because in electrolysis 
(see p. 55) it is attracted to the negative pole, and of the material 
combined with the hydrogen as the negative radical, because it is 
attracted to the positive pole. Thus the negative radicals in the 
above acids are CI, S0 4 , and C0 2 CH 3 , respectively. The first (CI) 
is a simple radical, the others are compound radicals. In many 
interactions the compound radicals move as units from one state 
of combination to another. 

Preparation by Displacement from Diluted Acids, — 

Ever}- one of the metals which displace hydrogen from water will 
also displace it from dilute acids. The acids must be diluted with 
water, unless, like hydrochloric acid, they are already dissolved in 
water. The action is much more vigorous than that on water, so 
that the most active metals are not employed. Metals like zinc, 
iron, and aluminium serve the purpose. The metal combines 
with the negative radical, and so liberates the hydrogen, which 
escapes in bubbles. Evaporation of the clear liquid, when the 
metal has all disappeared, gives in dry form the compound of the 
metal with the negative radical. Thus, with zinc and dilute 
sulphuric acid, zinc sulphate ZnS0 4 is produced. 

Skeleton: Zn + H 2 S0 4 -» ZnS0 4 + H. 

Balanced: Zn + H 2 S0 4 -> ZnS0 4 + 2H. 

With aluminium and hydrochloric acid, the product is aluminium 
chloride AICI3: 

Skeleton: Al + HC1 -> AICI3 + H. 

Balanced: Al + 3HC1 -> A1C1 3 + 3H. 



54 



COLLEGE CHEMISTRY 



The water undergoes no change during the action, although its 
presence is essential. It is simply a part of the apparatus. Any 
acid may be used, although with many the action goes on very 
slowly. 

For preparing small amounts of hydrogen, the apparatus (Fig. 
24) is such that additional acid may be added through the thistle-, 




Fig. 24. 



Fig. 25a. 



Fig. 25b. 



or safety tube. This avoids opening the flask and admitting air. 

The gas may be caught like oxygen over water or, being lighter 
than air, may be collected by downward displace- 
ment of the latter (Fig. 25a) . Heavy gases are 
collected by upward displacement of air (Fig. 
25b). 

With a Kipp's apparatus (Fig. 26) the gas 
may be made on a large scale and its delivery 
can be regulated. When the stream of gas is 
shut off by the stopcock, the pressure of the 
gas, as it continues to be generated, drives the 
acid away from the metal and up into the 
globe above, so that the action ceases. Yet 
the action is ready to begin again the moment 
any portion of the stored gas is drawn off for use. 
Silver, gold, and platinum, which do not 
combine with free oxygen, and even copper 
and mercury, which do, are all unable to lib- 
Fig. 26. erate hydrogen and to form oxides when heated 

in steam. When treated with dilute acids, none of these metals 

is able to displace and liberate the hydrogen (see order of activity 

of the metals, p. 59). 




HYDROGEN 



55 



Contact of the zinc or iron with an inactive metal, like platinum 
or copper, forms an electrical couple and hastens the interaction. 
For the same reason, commercial zinc, which contains traces of 
other metals, gives a steady evolution of hydrogen, while extremely 
pure zinc is almost inactive. 



The Third Variety of Chemical Change: Displacement. 

— The reactions used in liberating hydrogen illustrate the third 

of the four common forms of chemical change. Here a simple 

substance (the metal) and a compound (the acid) interact; the 

compound is divided into its radicals; 

and the simple substance combines 

with one radical while the other radical 

is liberated. The interacting element, 

here the metal, is said to displace the 

other element, here the hydrogen, from 

combination. The action of metals on 

water is a displacement also. 

Preparation of Hydrogen by 
Electrolysis. — If we dissolve any 
acid in water, and immerse the wires 
from a battery in the solution, bubbles 
of hydrogen begin to appear on the 
negative wire (the cathode) and rise to 
the surface . All the other constituents, 
whatever they may be, are attracted 
to the positive wire (the anode) and, 
therefore, do not interfere with the 
collection of pure hydrogen. An appa- 
ratus devised by Hofmann (Fig. 27) 
enables us to secure the hydrogen, 
which ascends on the left and accumu- 
lates at the top of the tube, displacing 
the solution. When hydrochloric acid is used: HC1 — > H (neg. wire) 
+ CI (pos. wire), the chlorine, a soluble gas, remains dissolved in 
the water near thepositive pole. When sulphuric acid is employed: 




H2SO4 — > 2H (neg. wire) + S0 4 (pos. wire), 



(1) 



56 COLLEGE CHEMISTRY 

The SO4, however, acts upon the water: 

S0 4 + H 2 -> H 2 S0 4 + 0. (2) 

Thus, the sulphuric acid is re-formed, round the positive wire, 
and only hydrogen and oxygen are finally liberated. 

Decomposition of a compound by the use of electrical energy is 
called electrolysis (Gk., decomposition by electricity). 

The Other Ways of Preparing Hydrogen. — For special 
purposes, hydrogen may be made by boiling an aqueous solution 
of sodium hydroxide with aluminium turnings, when sodium alumi- 
nate is formed: Al + NaOH + H 2 -> NaA10 2 + 3H; also by 
heating powdered zinc and dry sodium hydroxide, the product 
being sodium zincate: Zn + 2NaOH — > Na 2 Zn0 2 + 2H. 

Sources of the Hydrogen of Commerce. — Zinc is too 
expensive a substance to use in the preparation of hydrogen in 
large quantities for commercial purposes. We realize this when 
we note that 33 parts of zinc will liberate only one part of hydrogen, 
so that with 1 lb. of zinc we obtain only one half-ounce of the gas. 
Different sources are used in different localities and countries. 

The largest supply is probably obtained as a by-product in the 
electrolysis of an aqueous solution of common salt (NaCl), in 
connection with the manufacture of caustic soda (sodium hy- 
droxide, q.v.). The hydrogen is collected and compressed in steel 
cylinders. 

In some circumstances, the method of passing steam over heated 
iron is used (p. 51). 

Another plan is to liquefy water-gas (q.v.), a mixture of hydrogen 
and carbon monoxide. The hydrogen evaporates much the more 
readily of the two, and can thus be separated. This, and still 
other processes, involve substances and reactions which we have 
not yet encountered and will be mentioned at the appropriate points. 

Physical Properties. — Hydrogen is a colorless, tasteless, 
odorless gas. One liter weighs only 0.08987 g., while one liter of 
air weighs 1.293 g. Air is thus 14.5 times heavier, and hydrogen 
can be poured upwards (Fig. 28) and is used for filling balloons. 
Hydrogen was first liquefied in visible amounts by Dewar (1898). 



HYDROGEN 



57 



Ihe critical temperature is —234°. The colorless liquid boils at 
— 252.5° and, when allowed to evaporate rapidly under reduced 
pressure, freezes to a color- 
less solid (m.-p. -260°). All 
other gases, except helium, 
solidify easily when led into a 
vessel surrounded by liquid 
hydrogen. 

Hydrogen is even less sol- 
uble in water than is oxygen, 
1.8 volumes of the gas dissolve 
in 100 volumes of water at 15°. 
Hydrogen is absorbed, for the 
most part in a purely mechan- 
ical way, by many metals. 
Heated iron will take up 19 Fig. 28. 

times its volume of hydrogen, gold takes up 46 volumes, platinum 
in fine powder 50 volumes, palladium 502 volumes, and silver none. 
The maximum absorbed by palladium under favorable conditions 
is 873 volumes. 




Diffusion, — When two cylinders, one filled with hydrogen 
and one with air, are placed mouth to mouth (Fig. 29), so that the 
one containing hydrogen is uppermost, since the air 
in the lower cylinder is 14.5 times heavier than the 
hydrogen, we might expect the gases to remain in 
their respective cylinders. The air, however, makes 
its way into the hydrogen above it, and the hydrogen 
penetrates into the air in the lower cylinder so that, 
in a short time, the gases are perfectly mixed, just as 
if gravity did not exist. The same phenomenon is 
observed when, in everyday life, a bottle of scent is 
opened. The vapor, on escaping, begins to penetrate in all 
directions through the room, showing its presence by its odor. 
The material of gases has in fact an independent power of loco- 
motion. The resulting phenomenon we call diffusion. It is 
constant in rate for each gas under like conditions, and hydrogen 
has the greatest speed of diffusion of all the gases. 
The different rates of diffusion of different gases are easily 



Fig. 29. 



58 



COLLEGE CHEMISTRY 



shown by comparing their several speeds with that of air, when 
both pass through a wall of unglazed, porous porcelain. 

The porous cylinder A (Fig. 30) contains air and is connected by 
a rubber stopper with a wide tube which dips beneath the surface 
of the water. When a cylinder H containing hydro- 
gen is brought over it, rapid escape of gas takes place 
through the water, showing that a rise in pressure has 
taken place inside the porous vessel. Before the cylin- 
der of hydrogen approached the porous vessel, the air 
was moving both outwards and inwards through the 
porcelain, but, being the same air, the speed of motion 
was equal in both directions, and therefore the pres- 
sure inside was not affected. It is important to note 
that there was at no time rest, there was simply equal 
motion in both directions* When the hydrogen at- 
mosphere surrounded the cylinder, the hydrogen gas 
moved more rapidly into the cylinder than the air in- 
side could move out, and hence an excess of pressure 
quickly arose in the interior. 

Exact measurement shows that the lighter a gas is 
in bulk, the faster its parts move by diffusion in any 
direction. The rate is inversely proportional to the square root of 
the density of the gas. Thus, for hydrogen and air it is in the ratio 
VL293 : Vo.0897, or 3.8 : 1. 




Fig. 30. 



Fig. 31. 



^> 



Chemical Properties, — Hydrogen, 
delivered from a jet, burns in air or pure 
oxygen. A cold vessel held over the 
almost invisible blue flame condenses to 
droplets of water the steam that is pro- 
duced (Fig. 31). When hydrogen and 
oxygen are mingled in a suitable burner 
(Fig. 32), although the flame gives little 
light, it is exceedingly hot. Platinum 
melts in it easily and an iron wire burns brilliantly. In a closed 
space it produces a temperature of over 2500°. When the flame is 
allowed to play on a piece of quicklime, the latter becomes white- 
hot at the spot where the flame meets it. This result is called a 
calcium light or lime light. 




Fig. 32. 



HYDROGEN 59 

When hydrogen and oxygen are mixed, the chemical action is 
very slow at ordinary temperatures, no perceptible amount of union 
occurring in a period of five years. If the mixture is sealed up and 
kept at 300°, after several days a small part is found to have com- 
bined to form water. At 518°, hours are required before the union 
is complete. At 700° the combination is almost instantaneous. 
Hence contact with a body at a bright-red heat is required actually 
to explode the mixture. 

Finely divided platinum, when held in the cold mixture, hastens 
the union (otherwise vanishingly slow) in the part of the gases in 
contact with it. The heat of the union raises the temperature of 
the platinum and of neighboring portions of the gas and causes 
explosion of the mass. The platinum is simply a contact agent 
(p. 29) and remains itself unaffected. 

Hydrogen unites directly with a minority only of the simple 
substances. It combines rapidly with oxygen, chlorine, fluorine, 
and lithium, and more slowly with a few others. 

Hydrogen acts also upon some of the compounds of metals with 
oxygen or chlorine. Thus, when any one of the oxides of iron is 
heated in a tube through which hydrogen flows, the latter com- 
bines with the oxygen to form water, and the metal is liberated. 
The skeleton equation (p. 51) is: Fe 3 4 + H -> H 2 + Fe. We 
then reason that Fe 3 will give 3Fe. Since all the oxygen is removed 
from the compound, 4 will give 4H 2 0. To produce this, 8H is 
required. Hence: 

Fe 3 4 + 8H -> 4H 2 + 3Fe. 

This interaction is classed as a displacement. In describing it the 
chemist would also say that the hydrogen has been oxidized and 
that the oxide of the metal has been reduced (pp. 36-37). 

The Order of Activity of the Metals. — We employ metals 
so frequently in chemistry, that we must at once become familiar 
with the key to the main differences in their behavior. The 
order of their activity explains these differences, as well as many 
other facts. In the adjoining list, the most active metals are 
at the top. Hydrogen is not a metal, but is included because 
chemically it resembles the metals. All the metals above hydrogen 
displace this element from dilute acids (and from water), while 
those below it do not. 



00 



COLLEGE CHEMISTRY 



The first displaces the hydrogen from water violently, the second 
less vigorously. Magnesium barely acts on boiling water, but, like 
iron, acts on superheated steam. Zinc liberates hy- 
drogen with reasonable vigor from dilute acids, lead 
rather feebly, and copper and those following not at 
all. 

Other facts are explained by the table. Thus, when 
the metals are heated in pure oxygen, the last two 
do not combine. Those above silver do unite with 
oxygen — mercury rather slowly and the others more 
and more energetically as we ascend the list. Again, 
if we take the oxides of the metals, we find that those 
of the metals up to and including mercury lose all 
their oxygen when heated. If we heat the oxides, and 
lead hydrogen over them, the oxygen is easily removed 
from all the oxides up to and including those of iron, 
leaving in each case the metal. Thus, in general, 
the more active metals form the most stable com- 
pounds. 

The metals following hydrogen are the ones which 
are found in nature in large amounts in the free con- 
dition. 



Order of 

Activity. 
Metals 

Potassium 

Sodium] 

Calcium 

Magnesium 

Aluminium 

Manganese 

Zinc 

Chromium 

Iron 

Nickel 

Tin 

Lead 

Hydrogen 

Copper 

Bismuth 

Antimony 

Mercury 

Silver 

Platinum 

Gold 



Exercises. — 1. Make equations for reactions in which hydro- 
gen is liberated by the action of: (a) hydrochloric acid and mag- 
nesium giving MgCl 2 , (b) steam and zinc giving ZnO. 

2. Make an equation for the action of heat on manganese 
dioxide Mn0 2 giving oxygen and Mn 3 4 . 



CHAPTER VI 

VALENCE. CALCULATIONS 

Equivalence and Valence. — If the equations showing dis- 
placement of hydrogen by a metal be now re-examined, a peculiar- 
ity will be observed which we have thus far omitted to note. When 
sodium (p. 50) and calcium (p. 51) act upon water, one atomic 
weight (or atom) of the former displaces one atomic weight of 
hydrogen, but one atomic weight of the latter displaces twice as 
much hydrogen. Again, one atom of zinc (p. 53) displaces two 
atoms of hydrogen, but one atom of aluminium displaces three. 
Assuming, for simplicity, that we allow three of these metals all 
to act upon dilute hydrochloric acid, the equations are: 

Na + HC1 -> NaCl + H. 
Ca + 2HC1 -> CaCl 2 + 2H. 
Al + 3HC1 -> A1C1 3 +3H. 

Interpreting this, we perceive that the atom of aluminium, for 
example, displaces 3H, because it is able to combine with 3CI, and 
so incidentally liberates the hydrogen formerly united with 3C1. 
The atom of sodium, however, can unite with only 1C1, and so 
releases only 1H. Now this is not a rule confined to these re- 
actions, but represents a general chemical property of the atomic 
weight of each element, and a property which we shall find most 
useful. 

The atom of aluminium releases 3H because it can take the 
place of three atoms of hydrogen in chemical combination (and 
hold 3C1). The atomic weight of aluminium is said to be equiva- 
lent to (equal in chemical value to) three atomic weights of hydro- 
gen. Since it combines with 3 atomic weights of chlorine, it is also 
considered to be equi-valent to 3 atomic weights of this element. 

The chemical property referred to is called valence. The valence 
of an atomic weight of hydrogen or of chlorine is the unit. An 
atomic weight of sodium is said to be univalent, one of calcium 

61 



62 COLLEGE CHEMISTRY 

bivalent, one of aluminium trivalent. The formula H 2 shows the 
atomic weight of oxygen to be bivalent, because it unites with two 
atomic weights of hydrogen. Apparently, the atomic weight (or 
atom) of each element has a fixed capacity for combining with not 
more than a certain number of atomic weights (or atoms) of other 
elements. 

Marking the Valence. — Until we have become familiar with 
the valence of each element, it is advisable to mark the valences 
in a special way: Na 1 , Ca n , Al m , O 11 , Zn 11 , CI 1 . 

As we should expect, a bivalent atom can combine with two 
univalent atoms, or with one bivalent atom, and so forth. Thus 
we have the compounds of oxygen: Na 2 I 11 , Ca n O n , Al 2 ra 03 n , 
Zn n O n , CVO 11 . 

The rule is that the quantities of two elements which combine 
must have equal total combining capacities — i.e., identical total 
valence. Thus, Ca 11 has the valence two, and so does O 11 . Again, 
Al 2 m has a total valence of 2 X 3 ( = 6) and so has 3 n (3x2 = 6). 

Frequently the valence is marked by means of lines, the num- 
ber of lines pointing towards a symbol indicating the valence of 
the atom it represents : 

.CI , CI 

Na-Cl Ca Ca = Al-Cl = Al-0-Al = 

X C1 X C1 

Definition. — The valence of an element is a number repre- 
senting the capacity of its atomic weight to combine with, or dis- 
place, atomic weights of other elements, the unit of such capacity 
being that of one atomic weight of hydrogen or chlorine. 

Valence of Radicals. — What we have said applies to com- 
pounds of not more than two elements — so called binary 
compounds. We cannot with certainty tell the valences in a 
compound of three or more elements, like H 2 SC>4. But we have 
seen that the acids behave as if composed of two radicals: H(C1), 
H 2 (S04), that is, of two groups which move as wholes in chemical 
reactions. Hence we can assign a valence to a compound radical 
as a whole. Thus, (S0 4 ) n is evidently bivalent, as a whole, be- 
cause it is united with 2H 1 . Na(OH) and Ca(OH) 2 show the 
radical hydroxyl (OH) to be univalent. 



VALENCE. CALCULATIONS 63 

It is to preserve the identity of the radicals, and to make them 
easily recognizable, that we write them in brackets and place the 
coefficient outside, as Ca(OH) 2 and A1 2 (S0 4 )3, instead of using the 
forms Ca0 2 H 2 , Al^O^, and so forth. In fact, substances which 
commonly interact as if the radicals were single elements, we re- 
gard as binary compounds. 

In writing formulae of inorganic compounds we usually place 
the positive radical (p. 53) in front and the negative radical after it. 

Use in Making Formulse and Equations. — The chief use 
of the conception of valence is the very practical one of enabling 
us to write formulae. In making equations we constantly need to 
know whether the chloride of an element, say magnesium, is 
MgCl, or MgCl 2 , or MgCIs, or MgCLi, etc., and whether its sul- 
phate is MgS0 4 , or Mg 2 S0 4 , or some other combination of the 
symbols. To answer questions like this it is not necessary to 
know the formula of every compound of each element; the ap- 
parent disorder of these numbers can be reduced to rule, and the 
reader should endeavor thoroughly to master the rule before going 
farther. 

Thus, suppose that we require the formula of aluminium hy- 
droxide. Up to this point, we should have been compelled to 
look for it in a book. And if, later, we needed the formula of 
aluminium sulphate, we should have had to look that up, sepa- 
rately, also. But now, all we need is to know the valence of 
aluminium Al m , of the hydroxyl radical (OH) 1 and of the sul- 
phate radical (SO4) 11 . Making the total valences in the two 
halves of each compound alike, we write the formulae Al ni (OH) 3 I , 
Al 2 in (S0 4 )3 n . 

The reader must make a special effort to note the valences of 
each element and radical, and always to use them in making 
formulae. If a formula is written from memory, the valences 
must be checked, to make sure that the formula is correct. 

How to Learn the Valence of an Element. — To find out 
the valence of an element, we must obtain the formula of one simple 
compound of the element, containing another element of known 
valence. Thus, what is the valence of carbon? Its oxide is C0 2 . 
The total valence of oxygen here is 2 X 2 = 4. Carbon C IV is 



64 COLLEGE CHEMISTRY 

therefore quadrivalent. Hence its chloride must be C^CV (carbon 
tetrachloride), and it should give a compound with hydrogen C^W 
(methane, composing a large part of natural gas). When carbon 
combines with a trivalent element, equi-valent amounts of each 
element must be used, as in Al4 m C 3 IV (aluminium carbide), where 
AI4 111 and C 3 IV contain 3 X 4, or 12 units of valence each. 

The chemist does not memorize the valences themselves; he 
recovers the valence of an element or radical, when needed, by recall- 
ing the formula of a substance containing this element or radical 
in combination with a more familiar element or radical, such as CI 1 
orH 1 . 

Elements with More than One Valence, — The rule of 
valence is somewhat complicated by the fact that many elements 
show more than one valence. In other words, the combining 
capacity of an atomic weight of such an element may have two 
(or even more) values, according to the conditions under which 
the action takes place. 

Thus, we have encountered two chlorides of iron, ferrous chloride 
Fe II Cl2 I and ferric chloride Fe™^ 1 . We have, in fact, two com- 
plete series of compounds of iron, such as: 

Bivalent (Ferrous): FeCl 2 , FeO, FeS0 4 . 
Trivalent (Ferric): FeCl 3 , Fe 2 3 , Fe 2 (S0 4 ) 3 . 

When an element forms two such series of compounds, we always 
call particular attention to the fact. 

Exceptional Valences, — Some elements show an exceptional 
valence in one compound. The valences shown in series of com- 
pounds are the important ones, and the exceptions need not 
particularly concern us. Thus, in addition to the oxides FeO and 
Fe 2 3 , iron gives the magnetic oxide Fe 3 C>4, where the valence of 
iron appears not to be a whole number, but § or 2§. Hence the 
valence is made regular by supposing the oxide to be a compound 
of the other two oxides, as if the formula were Fe II 0,Fe 2 III 3 . 

Nomenclature. — The names of compounds containing only 
two elements (the true binary compounds) end in ide. Such are 
the oxides j as ferric oxide Fe20 3 ; the carbides, as aluminium car- 



VALENCE. CALCULATIONS 65 

bide AI4C3; the chlor ides, as sodium chloride NaCl; the sulphides, 
as ferrous sulphide FeS, etc. 

When an element forms two (or more) compounds with another 
element, they are frequently distinguished thus: carbon dioxide 
C0 2 , carbon monoxide CO; phosphorus pentoxide P 2 5 , phosphorus 
^rioxide P 2 3 . 

To distinguish two compounds of the same elements, another 
plan is also used: ferrcws chloride FeCl 2 , ferric chloride FeCl*; 
mercurows oxide Hg^O, mercuric oxide HgO. The suffix ous indi- 
cates that the metal is combined with the smaller proportion of 
the negative element, and ic that it is combined with the larger 
proportion. 

The tendency — although not a universal rule — is to use the 
latter plan with compounds containing a metal and the former 
with compounds containing only non-metals. 

Equivalent Weights. — In the foregoing discussion of valence, 
we have more than once used the word " equivalent." For 
example (p. 61), it was stated that the atomic weight of aluminium 
is equivalent to three atomic weights of hydrogen, because it dis- 
places them, and to three atomic weights of chlorine, because it 
combines with that number: 

Al + 3HC1 -> AICI3 + 3H 

Weights: 27.1 3 X 36.468 27.1 + 3 X 35.46 3 X 1.008 

Now chemists often view this from the other direction, and say 
that 1.008 g. of hydrogen are displaced by 9.03 g. of aluminium 
(one-third of the atomic weight) and that 35.46 g. of chlorine 
combine with only 9.03 g. of aluminium. When taking this view, 
they refer to the weight of an element displacing one atomic weight 
of hydrogen, or combining with one atomic weight of chlorine (or of 
any other univalent element) as the equivalent weight of that ele- 
ment. The equivalent weight of aluminium is therefore 9.03 and 
that of calcium Ca 11 20 (one-half the atomic weight) and that of 
sodium Na 1 23 (the atomic weight). 

It will be seen that the equivalent weight can always be found 
by a quantitative experiment. It is also evident that it is equal 
to the atomic weight divided by the valence. It is likewise clear 
that the equivalent weight of an element, multiplied by the valence 



66 COLLEGE CHEMISTRY 

of that element, is equal to the atomic weight. The conception 
of equivalent weights finds application in several connections in 
chemistry (see Normal Solutions and Faraday's Law). 

Calculations 

As we have seen (p. 44), the formula represents the composition 
of a substance, using the atomic weights as the units. We have 
learned how the formula is calculated from measurements made 
in an experiment (p. 45). We may now take up some of the ways 
of using the information contained in a formula. 

Composition from the Formula. Formula- Weight. — 

To learn the composition of a substance, such as potassium 
chlorate, KCIO3, from its formula, we look up the values of the 
atomic weights (inside rear cover). We find K = 39.1 parts of 
potassium, CI = 35.46 parts of chlorine, and 3 = 3 X 16.00 or 
48 parts of oxygen. The proportions, in order, are therefore: 
39.1 : 35.46 : 48. 

What is the proportion of oxygen to potassium and chlorine, 
together? It is 48 : 39.1 + 35.46, or 48 : 74.56, or 1 : 1.55. 

We require a name for the sum of the weights of the constituents 
indicated in the formula. This is called the formula-weight. 
Thus, for potassium chlorate, it is 39.1 + 35.46 + 48, or 122.56. 

To Find the Percentage Composition. — In potassium 
chlorate the proportions are 39.1 of potassium, 35.46 of chlorine, 
and 48 of oxygen or a total of 122.56. In one hundred parts, the 

potassium is ' X 100, or 31.9%; the chlorine ' fi X 100, 

48 
or 28.9%; and the oxygen t^Tq x 100 > or 39.1%. 

Stated in terms of the rule of proportion, we have, for the potas- 
sium, 122.56 : 39.1 :: 100 : x, where x is the percentage of potassium. 

Calculations by Use of Equations. — We frequently wish 
to know what weight of a product can be obtained from a given 
weight of the necessary materials. For example, what weight of 
ferrous sulphide can be made with 100 g. of iron? It is under- 
stood that the necessary sulphur is available. 



CALCULATIONS 67 

To avoid the blunders which are easily made, observe strictly 
the following rules: 

1. Write down the equation: 

Fe + S -> FeS. 

2. Place under each formula the weight it represents: 

Fe + S -> FeS 
55.84 32.06 87.90 

3. Read this expanded equation. In this case it reads: 55.84 
parts of iron combine with 32.06 parts of sulphur to give 87.90 
parts of ferrous sulphide. 

4. Re-read the original problem: "What weight of ferrous 
sulphide can be made with 100 g. of iron? " Having done this, 
place the amount given in the problem (100 g. of iron) under the 
formula of the substance in question. Then notice what the prob- 
lem asks ("what weight of ferrous sulphide") and place an x under 
the formula of that substance : 



Fe 


+ 


S 


— » 


FeS 


55.84 




32.06 




87.90 


100 g. 








X 



5. Read the problem as now tabulated: 55.84 g. of iron give 87.90 
g. of ferrous sulphide, therefore 100 g. of iron will give x g. of 
ferrous sulphide. 

6. State the proportion in this order (or, see below). 

55.84 : 87.90 :: 100 : x ( = 157.4 g). 

If the tabulation in rule 4 has been prepared correctly, this final 
statement as a proportion is purely mechanical. It will be noted 
that only two of the three quantities given in the expanded equa- 
tion were actually used. 

6a. Alternative Method. At the sixth step, we may also say: 
If 55.84 g. of iron give 87.90 g. of ferrous sulphide, 1 g. of iron will 

87.90 
give ' g. (= 1.574 g.) of ferrous sulphide. Then, if 1 g. of 

iron gives 1.574 g. of ferrous sulphide, the 100 g. of iron will give 
100 X 1.574 g. (= 157.4 g.) of ferrous sulphide. 



68 COLLEGE CHEMISTRY 

Warnings, — In solving the exercises at the end of the chap- 
ter, beware of three kinds of mistakes, which are commonly made. 

1. Do not read the problem carelessly and make the equation 
backwards, that is, with the sides reversed. Focus attention first 
on the exact chemical change involved. 

2. Do not speak, or think of the symbols Fe and S as standing 
for "1 part" of iron or sulphur. They stand for 1 chemical unit, 
or atomic weight, or atom, in each case, that is, for " 55.84 parts" 
and " 32.06 parts," respectively. 

3. Follow the rules laid down above. The chemist follows 
these rules. The beginner always thinks he can do without them, 
and he fails in consequence. Writing the equation in expanded 
form (rule 4) and reading the problem into it (rule 5) are abso- 
lutely essential steps. 

Another Example. — What weight of hydrogen is required 
to reduce 45 g. of magnetic oxide of iron to metallic iron? 

Following the rules, as before, we reach the expanded equation: 



Fe 3 4 


+ 


4H 2 


-> 3Fe 


+ 


4H 2 0. 


3 X 55.84 + 4 X 16 




8 X 1.008 


3 X 55.84 




4(2 X 1.008 + 16) 


167.52 + 64 




8.064 


167.52 




4 X 18.016 


231.52 




8.064 


167.52 




72.064 


45 g. 




X 









Observe that the atomic weights are multiplied by the sub- 
numbers, so that, for example, Fe3 = 3 X 55.84. Observe also 
that the formula-weights are multiplied by the coefficients, when 
such occur in front of the formulae, so that, for example, 4H 2 = 4 
X 18.016. 

The proportion 231.52 : 8.064 :: 45 : x (= 1.57) supplies the 
answer, 1.57 grams of hydrogen. 

Using the alternative plan: If 231.52 g. of magnetic oxide are 

reduced by 8.064 g. of hydrogen, 1 g. will be reduced by ' 2 g. 

(= 0.035 g.) of hydrogen. Hence, if 1 g. of magnetic oxide is 
reduced by 0.035 g. of hydrogen, 45 g. will be reduced by 
45 X 0.035 g. (= 1.57 g.) of hydrogen. 

Exercises. — 1. What are the valences of the negative radicals 
of phosphoric acid H3PO4, and of acetic acid (p. 52)? What must 



CALCULATIONS 69 

be the formulae of calcium phosphate, cupric acetate (Cu 11 ), alumin- 
ium phosphate, ferrous carbonate (C0 3 n ), ferrous sulphate, 
cupric chloride? 

2. What is the valence of phosphorus in phosphorus pentoxide 
(p. 32)? What must be the formulae of, (a) the corresponding 
chloride and sulphide of phosphorus, and (b) of aluminium oxide? 

3. What are the valences of the elements in the following: LiH, 
NH 3 , SeH 2 , BN? 

4. What are the valences of the metals and radicals in the fol- 
lowing: HN0 3 , Pb(N0 3 ) 2 , Ce(S0 4 ) 2 , KC1, KMn0 4 (potassium 
permanganate)? Name all the substances in 3 and 4. 

5. Make equations to represent, (a) the reduction of lead dioxide 
(Pb0 2 ) by hydrogen, (6) the actions of aluminium upon cold water 
and, (c) upon steam at a red heat. 

6. What weight of mercury is obtained from 120 g. of mercuric 
oxide HgO? 

7. What weight of mercuric oxide will furnish 20 g. of oxygen? 

8. What weight of Fe 2 3 may be obtained from 10 g. of oxygen? 

9. How much silver is contained in 100 g. of an impure specimen 
of silver chloride AgCl which is 33 per cent sand? 

10. What are the percentage compositions of cerium sulphate 
Ce(S0 4 ) 2 , phosphorus pentachloride PCI5, and ammonium chloride 
NH4CI? 

11. What weight of hydrogen is required to reduce 100 g. of 
ferric chloride FeCl 3 to ferrous chloride FeCl 2 ? 



CHAPTER VII 

THE MEASUREMENT OF QUANTITY IN GASES. RELATIONS 
BETWEEN STRUCTURE AND BEHAVIOR OF MATTER 

A specimen of a gas, like a specimen of a solid or a liquid, may be 
weighed, but it is usually easier to determine the quantity of the 
gas by (1) measuring its volume, and at the same time (2) noting 
its temperature on a thermometer suspended in it or close to it, and 
(3) ascertaining the pressure which it exercises. 

The Measurement of the Pressure of a Gas. — In almost all 
cases the easiest way to take account of the pressure of a gas is to 
place it in an apparatus so constructed that one 
boundary of the volume is a liquid. The apparatus 
is then so adjusted that the surface of the liquid in 
contact with the gas in the closed tube (Fig. 33) is at 
the same level as the free surface of the liquid which is 
exposed to the atmosphere. The equality in the levels 
of the liquids is then a guarantee that the specimen 
of gas and the atmosphere are exercising equal pres- 
sures on the liquid. At this stage the volume of 
the gas is measured, by reading the graduation (not 
shown) on the tube. Simultaneously the pressure of the atmos- 
phere and, therefore, of the gas, is ascertained by reading the 
barometer. 

The barometer (Fig. 34) consists of a bent tube containing mer- 
cury. The short limb (to the left) is open and the pressure of the 
atmosphere is exercised on the surface of the mercury there. The 
longer limb (to the right) is closed at the top and in it there is no 
gas above the mercury. When the tube is inclined, the surface of 
the mercury in the longer limb endeavors to retain the same verti- 
cal height above the lower surface and consequently rises and, with 
sufficient inclination, will reach entirely to the top of the tube. 
The downward pressure of the mercury on the right, above the 

70 





(^ 

















THE MEASUREMENT OF QUANTITY IN GASES 



71 



dotted line, is exactly equal to the pressure of the atmosphere on the 
free surface of the mercury at the same level. The amount of the 
latter pressure is proportional to the length of the column of mer- 
cury above the dotted line. Hence, reading the 
height at which the mercury stands above the free 
surface gives us a measure of the pressure of the 
atmosphere and of any specimen of gas which is 
at the same pressure. 

This is called the uncorrected reading. It is 
immediately reduced to the reading which would 
have been made if the barometer and its mercury 
had been at 0° {corrected reading), by noting the 
temperature on the adjacent thermometer and 
subtracting from the uncorrected reading the 
necessary correction (Table of Corrections, C, 
Fig. 34). 

For example: the volume of gas, after adjust- 
ment to atmospheric pressure, is 200 c.c. and its 
temperature 17°. The uncorrected barometric 
reading is 744 mm. with the barometer (perhaps 
in a different room from the gas) at 15°. The 
correction is —2.0 mm. The corrected reading is 
therefore 742 mm. 

Fig. 34. 

Correction of the Volume to 760 mm. Pressure. — 

Finally, since the atmospheric pressure varies from day to day, 
the volume at the observed pressure is corrected to that which the 
same quantity of gas would have occupied at the standard pressure 
of 760 mm. of mercury. By careful measurements, Boyle (1660) 
found that the volume occupied by the same sample of any gas 
varies inversely with the pressure. 

The illustration just given will show how this additional correc- 
tion is applied. There were 200 c.c. of the gas at 17° and 742 mm. 
(corr.). The question is: What would be the volume of this 
amount of gas at 760 mm.? At this new pressure (760 mm.) , which 
is greater than the old pressure (742 mm.), the volume will become 
less. Hence we change the volume in the proportion of these 
pressures, placing the smaller number in the numerator, so as to get a 
smaller volume as the answer: 200 X H£ = volume at 760 mm., 




72 COLLEGE CHEMISTRY 

= 195.3 c.c. If we wished to convert 100 c.c. at 775 mm. to 760 
mm., we should reason that the new pressure was smaller, and the 
volume would become greater, and should therefore place the 
larger number (775) in the numerator so as to get a larger volume 
for the answer. 

The Correction of the Volume of a Gas for Temperature. — 

The same sample of gas will occupy, when heated, a larger volume, 
and when cooled, a smaller volume than before. The change in 
volume for each degree Centigrade is ^t? of the volume of the same 
sample at 0°. To simplify the calculation we begin by converting 
the temperature to the absolute scale by adding 273° to each 
temperature. The volumes assumed by a sample of gas at different 
temperatures, the pressure remaining constant, are in the same pro- 
portion as the corresponding absolute temperatures (Charles, 1787). 
If the volume remains constant, then the pressure changes in the 
same proportion. 

In the illustration used above, there were 200 c.c. of gas at 17°, 
and it is required to know the volume at 0°. We add 278 algebra- 
ically to each temperature, and the question becomes: There are 
200 c.c. of gas at 290° Abs., what will be its volume at 273° Abs.? 
The volume changes in the direct ratio of the absolute temperatures. 
The new temperature is lower than the old, and the new volume will 
therefore be smaller than the old. Then 200 X § |f = volume at 
0° (273° Abs.) = 188.3 c.c. 

The above laws are usually applied to any example simulta- 
neously. Thus, 200 c.c. of gas at 742 mm. pressure (corr.) and 17° 
become 200 X iU X f|§ = 183.8 c.c. at 0° and 760 mm. 

Mixed Gases: Aqueous Tension, — Two gases at the same 
temperature, provided they do not interact chemically, do not in- 
terfere with each other's pressures when mixed. Thus, if they are 
forced into the same volume, the pressure of the mixture is equal 
to the sum of those of the components (Dalton's law, 1807). The 
gases are therefore still thought of individually, and the share 
which each gas has in the total pressure is called its partial pres- 
sure. This, like any other gaseous pressure, is proportional to 
the concentration of the particular gas in the mixture. 

For example, a gas measured over water contains water vapor. 



THE MEASUREMENT OF QUANTITY IN GASES 73 

The partial pressure of this, called aqueous tension (q.v.) } which 
is definite for each temperature, must be subtracted from the total 
pressure. The remainder is the partial pressure of the gas being 
measured, and this remainder is used as the pressure of this gas in 
any calculation. Thus, in a gas measured over water at 22°, the 
total pressure includes 19.7 mm. pressure of water vapor (the 
aqueous tension at 22°, see Appendix IV). Hence 150 c.c. of gas 
over water at 22° and 750 mm. is the same in amount as 150 c.c. of 
the same gas in dry condition at 22° and 730.3 mm. (there being 
simply 150 c.c. of water vapor at 19.7 mm. mixed with it). To 
obtain the volume of dry gas at 0° and 760 mm. we have the ex- 
pression 150 X tit X %f 3 . 

Densities of Gases.* — The density of a gas is the mass of 1 c.c. 
of the gas at 0° and 760 mm. pressure. Sometimes the weight 
of one liter (1000 c.c.) is called the density. Often the relative 
weight of the gas, the weight of an equal volume of air, or oxygen, 
or hydrogen being taken as unity, receives the same 
name. 

The most direct method of measuring the density 
of a gas is to employ a light flask of 125-150 c.c. 
capacity, provided with a rubber stopper and stop- 
cock (Fig. 35). By means of an air-pump the con- 
tents of the flask are removed, and it is weighed. 
This gives the weight of the empty vessel. The gas, 
whose density is to be ascertained, is then admitted, 
and care is taken that it finally fills the flask at 
the pressure of the atmosphere. The flask is FlG - 35 - 
closed and weighed again. The increase represents the weight of 
the gas. At the same time the temperature and barometric 
pressure are read. The volume is determined by displacing the 
gas once more from the flask, filling with water, and weighing 
again. The difference in weight between the empty flask and the 
flask full of water, in grams, represents the volume of the content 
of the flask in cubic centimeters. This volume is reduced to 0° 
and 760 mm. by the rules discussed above, and we have then a 
volume of the gas and the corresponding weight. 

* The subjects of this section are not actually used until Chapter IX (on 
Molar Weights) is reached. 




74 COLLEGE CHEMISTRY 

To illustrate, let us suppose that the volume of the flask is 200 c.c. 
and that it is filled with oxygen at 17° and 742 mm. The weight 
of the gas is found to be 0.26 g. We ascertained (p. 72) by calcula- 
tion that at 0° and 760 mm. this volume would be 183.8 c.c. The 
weight of a liter is given by the proportion 183.8 : 0.26 :: 1000 : x. 
Here x = 1.415 g. When the operation is performed carefully, 
and the weighing carried to the nearest milligram instead of the 
nearest centigram, a result more nearly approaching the accepted 
one (1.429) may easily be reached. 

To get the density of oxygen referred to hydrogen as unity, we 
must divide the answer by the weight of a liter of hydrogen 
(0.08987 g.). In the above case the quotient is 15.74. The ac- 
cepted value is 15.90. The density referred to air as unity is 
similarly obtained by dividing by 1.293, the weight of a liter of air 
at 0° and 760 mm. pressure. 

By using a modification of the flask just described, it is possible 
to ascertain the weights of known volumes of the vapors of liquids 
and solids. A temperature sufficiently high to vaporize the sub- 
stance must be employed. The volume is reduced by rule to 0° 
and 760 mm. and the density (in this case known as the vapor 
density) is calculated as before. The reduction to 0° and 760 mm. 
pressure by rule gives, of course, a fictitious result. The vapor 
would condense to the liquid form before 0° was reached, if the 
cooling were actually carried out. But the value for the density 
as it would be at 0° and 760 mm. has to be calculated to facilitate 
comparison with the corresponding values for other substances. 
The results have no physical significance, but are highly important 
to the chemist. 

Relations between the Structure and Behavior of 

Matter 

We have seen that matter is composed of minute particles 
called molecules. Just as we can thoroughly understand the be- 
havior of a watch or an automobile engine only if we know the 
details of its structure, and how the parts work, so we can under- 
stand the physical and chemical behavior of matter in masses 
only if we are familiar with its ultimate mechanism. Hence, we 
must now take up the structure of matter in its three states, the 



STRUCTURE AND BEHAVIOR OF MATTER 75 

gaseous, the liquid, and the solid. In doing this, we shall keep 
constantly in view the connection between the molecular relations 
and the general behavior of the matter. 

The Molecular Structure of Gases. — The most noticeable 
fact about gases is that they can be compressed to an enormous 
extent. Oxygen at 760 mm., for example, can be reduced by 
pressure to one two-hundredth of its volume, or even less. The 
compression does not affect the individual molecules, and there- 
fore does not diminish the volume actually occupied by the oxygen, 
but it crowds the molecules closer together and diminishes to one 
two-hundredth the space between them. Compressing a gas is, 
in fact, mainly compressing the empty space of which it chiefly 
consists. To understand what follows, the reader must keep 
constantly before him a mental image of a jar of gas as consisting 
of small particles separated by relatively wide, empty spaces. 
The molecules are in rapid motion and move in straight lines, ex- 
cepting when they strike one another or the walls of the vessel. 

The Properties of Gases. — Let us now note the more obvious 
qualities of gases, printing in italics the fact concerning a mass of 
gas and in black type the property of the molecules which accounts 
for the fact. 

The most remarkable thing about a gas, considering the loose- 
ness with which its material is packed, is the total absence in it of 
any tendency to settling or subsidence. Since the molecules cannot 
be at rest upon one another, as the great compressibility shows, we 
are driven to conclude that they are widely separated from one 
another, and that they occupy the space, otherwise a complete 
vacuum, by constantly moving about in all directions. But a 
moving aggregate of particles which does not even finally settle 
must be in perpetual motion. We must, therefore, believe the 
molecules to be wholly unlike particles of matter in having perfect 
elasticity, in consequence of which they undergo no loss of energy 
after a collision. They must continually strike the walls of the 
vessel and one another and rebound, yet without loss of motion. 
The fact that each gas is homogeneous, efforts to sift out lighter or 
heavier samples having failed, requires the supposition that all the 
molecules of a pure gas are closely alike. 



76 COLLEGE CHEMISTRY 

The diffusibility of gases is due to the motion of the molecules, 
and their permeability to the space available to receive molecules 
of another gas. These two modes of behavior involve no additional 
molecular properties. The word " diffusion' ' is often thought to 
mean the property of a given mass of gas in virtue of which another 
gas can mix with the given mass. This property is not diffusibility 
but permeability. It is the other gas, which makes its way into 
the given gas, which is diffusing. Diffusion is spontaneous 
motion of the parts of a gas away from their original location. 
Unless this motion is into an empty space, the diffusing molecules 
must, of course, move into another body of gas. In the case of 
the jars of hydrogen and air (p. 57), each gas moved in part out of 
its original jar (diffused), and each received parts of the other gas 
into its jar (was permeated). 

Boyle 9 s Law and Charles 9 Law. — Passing now to Boyle's 
law (p. 71), the thing to be accounted for is that when a sample of 
a gas diminishes in volume, its pressure increases in the same pro- 
portion. Let the diagram (Fig. 36) represent a cylinder with a 
movable piston, upon which weights may be placed to resist the 
pressure. Now the pressure exercised by the gas under 
the piston cannot be like the pressure of the hand upon 
a table, since we have just assumed that the particles 
are not even approximately at rest, and the spaces 
between them are enormous compared with the size 
I of the molecules themselves. The gaseous pressure 

-J 1 must therefore be attributed to the colossal hailstorm 

which their innumerable impacts upon the piston 
produce. If this is the case, the compressing of a gas 
must consist simply in moving the partition down- 
wards, so that the particles as they fly about are gradu- 
ally restricted to a smaller and smaller space. Their paths become 
on an average shorter and shorter. Their impacts upon the walls 
become more and more frequent. So the pressure which this 
causes becomes greater and greater, and is proportional to the 
degree of crowding (the concentration) of the molecules. 

There are two other points to be added. When we diminish the 
volume to one-half, we find from experience that the pressure be- 
comes exactly, or almost exactly, twice as great. This must mean 



STRUCTURE AND BEHAVIOR OF MATTER 77 

that, although the particles are becoming crowded, they do not 
interfere with one another's motion, excepting of course where 
actual collision causes a rebound. Only in the absence of inter- 
ference would doubling the number of molecules per unit of volume 
give exactly double the number of impacts on the walls. Hence 
the molecules must have practically no tendency to cohesion. 
Finally, the molecules must be supposed to move in straight lines 
between collisions. 

A baseless idea, that the molecules of a gas repel one another, 
still lingers in some quarters. There is no evidence of this. The 
molecules pay almost no attention to one another, either by 
attraction or repulsion. 

Boyle's law therefore adds four more details concerning molec- 
ular behavior, namely, that the impacts of the particles produce 
the pressure, that the crowding of the molecules represents the con- 
centration of the material and that the particles move in straight lines 
and show almost no cohesion, since pressure and concentration are 
very closely proportional to one another. 

How, now, can we account for Charles' law (p. 72), according to 
which an increase in pressure (or in volume) results from heating 
a mass of rapidly moving molecules? The action of a particle 
colliding with a surface is measured in physics in terms of its 
mass and its velocity. It is evident that heating a cloud of mole- 
cules would not increase the mass of each, and it must therefore 
increase the velocity of each, since the kinetic energy of all becomes 
greater. 

Avogadro's Law. — The identical general behavior of all 
kinds of gases suggests that their structures may be all alike. 
Avogadro (1811), the professor of physics in Turin, put forward 
the hypothesis that the numbers of molecules in equal volumes of 
different gases, at the same temperature and pressure, might be 
equal. A more strict study of the properties we have been con- 
sidering, and of some additional facts, has since shown that no 
other conjecture than Avogadro's would be consistent with them. 
Thus it is now accepted as a fact, and is known as Avogadro's law. 
It may also be put in the form: At the same temperature and 
pressure, the molecular concentration of all kinds of gases has the 
same value. 



78 COLLEGE CHEMISTBY 

Diffusion, — The law of diffusion (p. 58) harmonizes with the 
conceptions of molecular structure without further additions to the 
latter. The speed of the hydrogen molecule at room temperature 
is 1840 meters per second. The masses of the oxygen and hydro- 
gen molecules are as 16 : 1, and the speeds of diffusion (p. 58) as 
Vl : VT6, or 1 : 4. Hence the speed of the oxygen molecule is 
one-fourth of 1840, or 460 m. per sec. 

Calculation shows the activity of the molecules to be such that, 
in air, the number striking a single square centimeter of surface 
per second would fill no less than twenty liters. 

Liquefaction of Gases: Critical Temperature. — Finally, 
gases can be liquefied by sufficient cooling and compression. This 
fact compels us to suppose that, after all, even gaseous molecules 
have a tendency to cohesion. This cohesion is scarcely perceptible 
so long as the gas is warm and is diffuse. Thus, 2 liters of oxygen at 
1 atmosphere pressure, when subjected to 2 atmospheres pressure, 
give 0.9991 liters instead of 1 liter. The additional contraction 
of 0.0009 liters (0.9 c.c.) is due to the effect of cohesion when 
the molecules are thus crowded closer together. The gases which 
are more easily liquefied than is oxygen show greater effects. 
Thus, 2 liters of sulphur dioxide at 760 mm., when subjected to 2 
atmospheres pressure, give only 0.974 liters, showing a contraction 
due to cohesion of 26 c.c. These data refer to 0°. At lower 
temperatures the contractions due to cohesion become rapidly 
greater. This cohesion is not of the nature of gravitational 
attraction. 

We can readily understand, therefore, that when the kinetic 
energy of the molecules is sufficiently reduced by cooling (namely, 
to, or below the critical temperature, see below), and the molecules 
are brought sufficiently close together, the tendency of the mdle- 
cules to cohere causes the gas to condense and assume the liquid 
form. In 1869 Andrews found that carbon dioxide could be 
liquefied at 0° by 38 atmospheres pressure, and at 30° by 71 atmos- 
pheres, but that above 31.35° it could not be liquefied by any 
pressure. He discovered that each gas has a critical temperature, 
as he called it. For carbon dioxide, this temperature can be ob- 
served by placing a heavy-walled, glass tube (Fig. 37), half-filled 
with liquid carbon dioxide, in a beaker of water, and gradually 



STRUCTURE AND BEHAVIOR OF MATTER 79 

raising the temperature of the latter. At 31.35°, the surface 
between the liquid and gas becomes hazy and vanishes. When 
the temperature falls once more, the surface re-appears at 31.35°. 
This shows that, with Faraday's " permanent" gases, a tempera- 
ture below the critical point had not been employed. 

The critical temperature of oxygen is —118°, of 
hydrogen —23-4°, of carbon dioxide 31.35°, of sulphur 
dioxide 156°, of water 358°. 

Another Deviation from the Laws of Gases. 
A Perfect Gas. — It may be added that when a gas 
is already under very high pressure, and very closely 
packed, an increase in the pressure does not produce quite 
as great a diminution in volume as Boyle's law leads us to 
expect. This reminds us that we are diminishing only Fia. 37. 
the space between the molecules, and not the volumes of the mole- 
cules themselves, and therefore not the total volume of the gas. 
When, on severe compression, the volume occupied by the mole- 
cules themselves has become an appreciable fraction of the whole 
volume, additional compression does not affect the whole volume, 
and the contraction is smaller than Boyle's law would indicate. 
Thus, 2 liters of hydrogen, even at one atmosphere pressure, when 
subjected to two atmospheres pressure, give 1.0006 liters, instead of 
1 liter. 

The last two effects (namely, those due to the tendency to 
cohesion of, and the space occupied by the molecules) are called 
deviations from the laws of gases. In consequence of these individ- 
ual deviations, there are not exactly equal numbers of molecules in 
equal volumes of any two different gases, at the same temperature 
and pressure. An imaginary gas, which exhibits neither deviation, 
called a perfect gas, is often referred to in discussing the behavior 
of gases. 

Summary. — We may now summarize the principal facts 
about gases in mass, appearing in italics above, with the corre- 
sponding features of the molecular relations, in heavy type, which 
we have added one by one. 



80 



COLLEGE CHEMISTRY 



Facts About Gases in Mass. 


Corresponding Relations of Molecules. 


Compressibility .... 


Vacuum + molecules widely separated. 


Diffusibility. . 










Molecules in rapid motion. 


Permeability . 










Empty space relatively large. 


Non-settling . 










Molecules perfectly elastic. 


Homogeneity . 










Molecules of any one substance closely alike. 


Pressure . . . 










Due to impacts of molecules. 


Boyle's law . . 










Pressure proportional to concentration of the 




molecules. Molecules move in straight lines 




and, when widely scattered, show no tendency 




to cohesion or to repulsion. 


Charles' law. ...... 


Rise in temperature increases the velocity, and 




therefore the kinetic energy of the molecules. 


Above and other facts . 


Avogadro's law. 


Law of diffusion . . . 




Gases can be liquefied . 


Molecules do possess a tendency to cohesion, 




which becomes conspicuous when they are 




cooled and closely crowded together. 



History of the Kinetic Molecular Theory. — This theory- 
was first suggested by Daniel Bernoulli (1738), who explained by 
its means the pressure and compressibility of gases. Lomonossov 
(1748) developed the theory very completely and by means of it 
explained Boyle's law and the effects of changes in temperature. 
He also anticipated from the theory the existence of the second 
deviation from the law of gases (1749), a discovery usually cred- 
ited to Dupre (1864). He likewise pointed out that there was 
no limit to the maximum velocity of a molecule, and therefore no 
upper limit of temperature, but that there must be a lower limit 
(the absolute zero) at which the molecules would be at rest (1744). 
This work was entirely forgotten, until attention was called to it 
in 1904 by Menschutkin. 

Similar views were expressed by Waterston (1845), but were 
still so much ahead of the time that the committee of the Royal 
Society did not approve the paper for publication, and it was dis- 
covered in the archives of the society, long afterwards, by Lord 
Rayleigh. The development of the theory, so far as it applies to 
heat, is therefore credited to Joule (1855-60) and, in respect to all 
properties of gases, to Kronig (1856) and Clausius (1857), who 
knew nothing of the earlier work. 



STRUCTURE AND BEHAVIOR OF MATTER 81 

Molecular Relations in Liquids, — The fact that even great 
pressures produce little diminution in the volume of a liquid shows 
that the free space, present in gases, is absent in liquids. The 
measured effects of various pressures show, for example in the case 
of water, that to reduce the volume to one-half would require, not 
doubling the pressure as in a gas, but increasing it from 1 to 10,000 
atmospheres. The molecules of a liquid are actually in contact 
with one another. 

The phenomena connected with surface tension, such as co- 
herence into drops, show that cohesion plays a much larger part 
in liquids than in gases. On the other hand, liquids which are 
capable of mixing (e.g., alcohol and water), when placed above one 
another in the same vessel, do mix, slowly, by diffusion. This 
indicates that motion of the molecules, although much impeded 
by friction, has not been annihilated by cohesion. The escape of 
vapor — that is, of part of the liquid in gaseous form — likewise 
proves that the molecules in the liquid are in motion. The rela- 
tions of liquid and vapor can be discussed most effectively in the 
next chapter, in connection with the case of water and steam. 

Molecular Relations in Solids. — The properties of solids 
differ from those of liquids chiefly in the fact that the solid has 
a definite form of which it can be deprived only with difficulty. 
This we may explain in accordance with the kinetic hypothesis by 
the supposition that the cohesion in solids is very much more 
prominent than in liquids. We obtain solids from liquids by 
cooling them; in other words, by diminishing the kinetic energy 
and therefore the velocity of the particles. The cohesive tendency 
of the latter is thus able to make itself felt to a greater extent. If, 
conversely, we heat a solid, or, according to the hypothesis, if we 
increase the speed with which the particles move, the body first 
melts and gives a liquid, and this finally boils and becomes a gas. 
The intrinsic cohesion of the particular substance can undergo no 
change, but the increasing kinetic energy of the particles steadily 
and continuously obliterates its effects. Yet some motion still 
survives in a solid. Thus we find that when the layer of silver is 
stripped from a very old piece of electroplate, the presence of this 
metal in the German silver or copper basis of the article is easily 
demonstrated. 



82 COLLEGE CHEMISTRY 

The tendency of all solids to assume crystalline forms, which 
show definite cleavage and other evidences of structure, distin- 
guishes them sharply from liquids. The force of cohesion in 
liquids is exercised equally in different directions. In solids it 
must differ in different directions in order that structure may re- 
sult. Since each substance shows an individual structure of its 
own, these directive forces must have special values in magnitude 
and direction in each substance. 

Crystallization, — A crystal arises by growth. When the 
process is watched, as it occurs in a melted solid or an evaporating 
solution, the slow and systematic addition of the material in lines 
and layers, as if according to a regular design, is one of the most 
beautiful and interesting of natural phenomena. The fern-like 
patterns produced by ice on a window-pane show the general 
appearance characteristic of crystallization in a thin layer. A 
larger mass in a deep vessel gives forms which are geometrically 
more perfect. From its very incipiency the crystal has the same 
form as when, later, its outlines can be distinguished by the eye. 
Hence the outward form is only an expression of a specific internal 
structure which the continual reproduction of the same outward 
form on a larger and larger scale leaves as a memorial of itself in 
the interior. 

Crystal Forms. — Crystalline form is continually used in 
identifying (pp. 2, 12, 19) the substances produced in chemical 
actions. The classification of crystalline forms is carried out 
according to the degree of symmetry of the crystals: 

1. Regular system. 5. Monosymmetric, or 

2. Square prismatic system. monoclinic system. 

3. Hexagonal system. 6. Asymmetric, or 

4. Rhombic system. triclinic system. 

The regular system presents the most symmetrical figures of all. 
Some forms which commonly occur are the octahedron (Fig. 38) 
shown by alum, the cube (Fig. 39) affected by common salt, and 
the dodecahedron (Fig. 40) frequently assumed by the garnet. 

The square prismatic system includes less symmetrical forms than 
the previous one, since the crystals are lengthened in one direction. 



STRUCTURE AND BEHAVIOR OF MATTER 



83 



Fig. 41 shows the condition in which zircon (ZrSiOJ, which fur- 
nishes us with the basis of certain incandescent illuminating 
arrangements, occurs in nature. The form of ordinary hydrated 
nickel sulphate (NiS0 4 ,6H 2 0) is similar to this. 





^ 



Fig. 38. 



Fig. 39. 



Fig. 40. 



V 



Fig. 41. 



The hexagonal system, like the preceding, frequently exhibits 
elongated prismatic forms, but the section of the crystals is a hexa- 
gon, instead of a square, and the termination is a six-sided pyramid. 
Quartz (Fig. 42), or rock crystal, is the most familiar mineral in 



j\ 



v 



Fig. 42. 




Fig. 43. 



Fig. 44. 



Fig. 45. 



this system. Calcite (C&COz), which is chemically identical with 
chalk, or marble, takes forms known as the scalenohedron (Fig. 43) 
and rhombohedron (Fig. 44), which are classified in a subdivision 
of this system. Indeed, recently it has become common to erect 



^o> 




& 



Fig. 46. 



Fig. 47. 



Fig. 48. 



this into a separate system (the trigonal), in which both quartz and 
calcite are included. 

The rhombic system includes the natural forms of the topaz, 
and of sulphur (Fig. 7, p. 12), as well as that of potassium perman- 
ganate (Fig. 45), potassium nitrate (Fig. 46), and many other 



84 COLLEGE CHEMISTRY 

substances. These crystals exhibit a good deal of symmetry, but 
their section is always rhombic, and hence the name. 

The monosymmetric system exhibits forms which have but one 
plane of symmetry. Gypsum (Fig. 47), which is hydrated calcium 
sulphate CaS04,2H 2 0, and feldspar (Fig. 3, p. 2) are minerals pos- 
sessing forms of this kind. Tartaric acid, rock candy (Fig. 48), 
potassium chlorate, and hydrated sodium carbonate (washing 
soda) belong to this system. 

The asymmetric system includes forms which have no plane of 
symmetry whatever. Blue vitriol (Fig. 52, p. 95), CuS04,5H 2 0, is 
one of the most familiar substances of this kind. 

Exercises. — The text cannot be understood unless some 
problems involving the laws of gases are actually worked. 

1. Reduce 189 c.c. of gas at 15° and 750 mm. to 0° and 760 mm. 

2. Reduce 110 c.c. of gas at — 5° and 741 mm. to 0° and 760 mm. 

3. Convert 500 c.c. of gas at 25° and 700 mm. to 18° and 745 mm. 

4. Reduce 250 c.c. of gas (standing over water) at 22° and 755 
mm. to the dry condition and to 0° and 760 mm. 

5. The density of a substance referred to air is 3.2. What is the 
density referred to hydrogen? What will be the volume occupied 
by 10 g. of the substance at 20° and 752 mm.? 

6. Describe two ways of obtaining crystals of a substance. 



CHAPTER VIII 

WATER 

Water is as necessary to life as is oxygen. The human body 
is saturated with it and, to make up for evaporation, as well as to 
aid in digestion and other life processes, it is a necessary part of 
our food. The ocean covers three-fourths of the earth's surface, 
and the "dry" land is, fortunately, far from being really dry. 

Physical Properties of Water. — A deep layer of water has a 
blue or greenish-blue color. At a pressure of 760 mm., water ex- 
ists as a liquid between 0° and 100°. Below 0° it becomes solid, 
above 100° a gas. Of all chemical substances it is the one which we 
use most, so that its physical properties, discussed below, should 
be studied attentively. Then, too, what is said of water is in 
general true of all other liquids, from which it differs only in 
details. 

The quantity of heat required to raise one gram of water one 
degree in temperature, at 15°, is called a calorie, the unit quantity 
of heat. The specific heat of any substance being the quantity 
of heat required to raise the temperature of one gram of the sub- 
stance one degree, the specific heat of water is 1. The values for 
other substances are all smaller (e.g., limestone 0.2). Thus the 
temperature of large masses of water, such as lakes and seas, 
changes more slowly, and within a smaller range, than that of the 
rocks and soil composing the land. The more constant tempera- 
ture of the water tends to regulate that of the air, and hence the 
climate of an island is less variable from season to season than is 
that of a continent. 

Ice. — The raising or lowering of the temperature of a gram of 
water through one degree involves the addition or removal of one 
calorie of heat. The conversion, however, of a gram of water at 
0° to a gram of ice at 0° requires the removal of 79 calories. The 

85 



86 COLLEGE CHEMISTRY 

mere melting of a gram of ice causes an absorption of heat to the 
same amount, called the heat of fusion of ice. At 0° a mixture of 
ice and water will remain in unchanged proportions indefinitely. 
Any cause which tends permanently to lower or raise the tempera- 
ture by a fraction of a degree, however, will bring about the disap- 
pearance of the water or of the ice, respectively. This temperature 
is called the melting- or the freezing-point. 

Water can be cooled below 0° (supercooled) without beginning 
to freeze, unless it is stirred, or " inoculated" by the addition of a 
piece of ice. Hence, the freezing-point is not defined as the point 
at which ice begins to form, for that point varies, and is always 
below 0°, but as the temperature of a mixture of ice and water. 

Steam. — At atmospheric pressure, water passes into steam 
rapidly at 100°, but at lower temperatures, and even when frozen, 
it does the same thing more slowly. It changes into steam, how- 
ever, only when the necessary supply of heat is forthcoming. 
One gram of water at 100°, in turning into a gram of steam at 100°, 
takes up 540 calories. This is called its heat of vaporization. 
Steam, in fact, contains much more internal energy than an equal 
weight of water at the same temperature, just as water, in turn, 
contains more energy than ice. 

Steam is a colorless, invisible gas. The visible cloud of fog, 
seen when steam escapes into cold air, is composed of minute 
drops of water, formed by condensation, and visible because they 
have surfaces and reflect light. 

The States of Matter: Transition Points, — Most sub- 
stances are known in three different states of aggregation, solid 
(crystalline), liquid, and gaseous. There is no magic about the 
number, three, however. Thus, sulphur has a vapor state, two 
liquid states, and several solid forms. There are even five forms 
of ice, and most solids probably exist in several different states. 

All transitions from one state to another take place at some 
definite temperature (when the pressure is fixed). Such temper- 
atures, when referring to the change from the solid to the liquid, 
and from the liquid to the gaseous state are called the melting- 
point or freezing-point, and the boiling-point, respectively, or in 
general, are known as transition points. 



WATER 



87 



Aqueous Tension and Vapor Pressure, — The quantity 
of the vapor present is defined by the gaseous pressure it exercises, 
the value being called the vapor pressure of water vapor (or of 
the vapor of any other volatile substance) in the location in 
question. 

The most significant fact about vapor pressure is that, when ex- 
cess of the liquid is present, the pressure of the vapor quickly reaches 
a definite maximum value for each temperature. In the absence 
of excess of the water, less than this maximum pressure may exist. 
More than the maximum pressure proper to a given temperature, if 
produced by compression, cannot be maintained, however, for a 
part of the vapor condenses to the liquid state. The magnitude 
of this maximum vapor pressure, at a given temperature, depends 
on the ability of the particular liquid to generate vapor. This max- 
imum vapor pressure is held, therefore, to represent the vapor 
tension of the liquid, at the given temperature, and this is a specific 
property of the substance. 

The vapor tension may be shown by allowing a few drops of 
water to ascend into a barometric vacuum (Fig. 49). The tube 
on the left shows the mercury when nothing presses 
on its surface. The tube on the right shows the 
result of admitting the water. The difference in 
the height of the two columns gives the value of the 
vapor pressure of the water vapor. With excess of 
water, the value is that of the vapor tension, called, 
in the case of water, the aqueous tension. 

The jacket surrounding the tube on the right 
enables us, by adding ice or warm water, to main- 
tain any temperature between 0° and 100°. When 
ice is used outside, and a piece of it is introduced 
into the vacuum, the vapor it gives off quickly 
reaches a pressure of 4.5 mm. The vapor pressure 
of the ice takes the place of 4.5 mm. of mercury in 
balancing the atmospheric pressure, and so the mer- 
cury column falls by this amount. Similarly, water 
at 10° causes a fall of 9.1 mm. and at 20° of 17.4 mm., so that these 
represent the mercury-height values of the aqueous tension at 
these temperatures. The quantity of water used makes no differ- 
ence, so long as a little more is present than is required to fill the 



1 



Fig. 49. 



88 COLLEGE CHEMISTRY 

available space with vapor. With ether, instead of water, at 
10°, the fall is 28.7 mm. 

With water at higher temperatures the fall of the mercury col- 
umn becomes much greater. At 50° it is 92 mm., at 70° it is 233.3 
mm., at 90° it is 525.5 mm., and at 100° it is 760 mm., or one at- 
mosphere. At 121° the aqueous tension is two atmospheres, at 
180° it is ten atmospheres (see Appendix IV). 

When water at a certain temperature has given the full amount of 
water vapor to the space above it that its aqueous tension permits, 
we say that the space is saturated with vapor. That concentration 
of vapor which constitutes saturation varies with the temperature 
of the water and depends, therefore, solely on the power of the water 
to give off vapor. It has nothing to do with the size of the space, 
and is even independent of other gases the space may already con- 
tain. Thus, if a little air is first placed above the dry mercury 
(Fig. 49), causing it to fall, the additional depression produced by 
adding water is the same as if the air had been absent (p. 72. See 
footnote to p. 11). 

Water Vapor in the Air, — The space immediately above the 
surface of the ground, which is mainly occupied by atmospheric 
air, is, on an average, less than two-thirds saturated with water 
vapor. That is to say, such air, when enclosed in a vessel con- 
taining water, will take up about one-half more than it already 
contains. The vapor of water at 100° in an open vessel displaces 
the air entirely and, if the required heat is furnished, the liquid 
boils. 

All our substances and apparatus have traces of water, derived 
from the atmosphere, condensed on their surfaces. This water is, 
in a sense, in an abnormal condition, for it does not evaporate even 
in dry air. It is observed to pass off in vapor, however, when we 
have occasion to heat the substance or apparatus. 

Molecular Relations of Liquid and Vapor. — When the 
water was introduced above the barometric column, the vapor, or 
gaseous water, could have resulted only from the spontaneous 
motion of the molecules in the liquid. Some of the molecules, 
moving near the surface, went off into the space above the water 
and became gaseous. To be consistent, we must also conclude 



WATER 89 

that the vapor above the water is not composed of the same set of 
molecules one minute as it was during the preceding minute. 
Their motions must cause many of them to plunge into the liquid, 
while others emerge and take their places. When the water is first 
introduced, there are no molecules of vapor in the space at all, so 
that emission from the water predominates. The pressure of the 
vapor increases as the concentration of the molecules of vapor 
becomes greater, hence the mercury column falls steadily. At the 
same time the number of gaseous molecules plunging into the 
water per second must increase in proportion to the degree to 
which they are crowded in the vapor. The rate at which mole- 
cules return to the water thus begins at zero, and increases steadily; 
the rate at which molecules leave the water maintains a constant 
value. Hence the rate at which vapor molecules enter the water 
must eventually equal that at which other water molecules leave 
the liquid. At this point, occasion for visible changes ceases and 
the mercury comes to rest. We are bound to think, however, 
of the exchange as still going on, since nothing has occurred to 
stop it. The condition is not one of rest but of rapid and equal 
exchange. Such, described in terms of molecules, is the state of 
affairs which is characteristic of a condition of equilibrium. The 
condition is dynamic, and not static. 

Equilibrium. — This term is used so often in chemistry, and is 
used in so unfamiliar a sense, that the reader should consider 
attentively what it implies. Three things are characteristic of a 
state of equilibrium: 

1. There are always two opposing tendencies which, when equi- 
librium is reached, balance each other. In the foregoing instance, 
one of these is the hail of molecules leaving the liquid, which is 
constant throughout the experiment. It represents the vapor ten- 
sion of the liquid. The other is the hail of returning molecules, 
which, at first, increases steadily as the concentration of the vapor 
becomes greater. This is the vapor pressure of the vapor. These 
have the effect of opposing pressures and, when the latter becomes 
equal to the former, equilibrium is established. In all cases of 
equilibrium we shall symbolize the two opposing tendencies by two 
arrows, thus: 

Water (liq.) ^ Water (vapor). 



90 COLLEGE CHEMISTRY 

2. Although their effects thus neutralize each other at equilib- 
rium, both tendencies are still in full operation. In the case in point, 
the opposing hails of molecules are still at work, but neither can 
effect any visible change in the system. Equilibrium is a state, 
not of rest, but of balanced activities. 

3 (and this is the chief mark of equilibrium). A slight change in 
the conditions produces, never a great or sharp change, but always, 
and instantly, a corresponding small change in the state of the 
system. The change in the conditions accomplishes this by favoring 
or disfavoring one of the two opposing tendencies. Thus, for ex- 
ample, when the temperature of a liquid is raised, the kinetic energy 
of its molecules is increased, the rate at which they leave its surface 
becomes greater, the vapor tension increases and, hence, a greater 
concentration of vapor can be maintained. The system, therefore, 
quickly reaches a new state of equilibrium in which a higher vapor 
pressure exists. A heap of matter on a table is not in equilib- 
rium, because addition of more material produces no response 
until, when a very great quantity is added, the table breaks. 
But a body on the scales is in equilibrium, for the addition of 
the smallest particle produces a corresponding inclination of the 
beam. 

In the preceding illustration, the evaporating tendency was 
favored by a rise in temperature. As an example of a change in 
conditions disfavoring one tendency, take the case where the liquid 
is placed in an open, shallow vessel. Here the condensing tendency 
is markedly discouraged, for there is practically no return of the 
emitted molecules. Hence complete evaporation takes place. Ele- 
vation of the temperature hastens the process. A draft insures the 
total prevention of all returns, and has therefore the same effect. 
The two methods of assisting the displacement of an equilibrium, 
and particularly the second, in which the opposed process is weak- 
ened and the forward process triumphs solely on this account, 
should be noted carefully. They are applied with surprising effec- 
tiveness in the explanation of chemical phenomena (see Chaps. XIV 
and XVIII). 

Water as a Solvent, — One of those physical properties of 
water which are most used in chemical work is its tendency to dis- 
solve many substances. This subject is so important and exten- 



WATER 91 

sive that we shall presently devote a complete chapter to some of 
its simpler and more familiar aspects. 

Natural Waters. — The foreign material in natural waters is 
divided into dissolved matter and suspended matter. Rain-water, 
collected after most of the dust has been carried down, is the 
purest natural water. It contains, however, nitrogen, oxygen, and 
carbon dioxide dissolved from the air. Sea- water holds about 3.6 
per cent of dissolved material. River and, especially, well waters 
dissolve various substances during their progress over or under 
the surface of the ground. They often contain calcium sulphate, 
calcium bicarbonate, and compounds of magnesium, and are then 
described as hard. Sometimes they contain compounds of iron, 
and sometimes they are effervescent and give off carbon dioxide. 
These are called mineral waters. 

Many river waters contain large amounts of clay and organic 
matter (often due to admixture of sewage) suspended in them. 
It is not the organic matter which is deleterious, but the bacteria 
of putrefaction and disease which are present also, and are usually 
for the most part attached to the particles of suspended matter. 
Cholera and typhoid fever are often spread by the drinking of 
water into which sewage, infected by other patients suffering from 
these diseases, has been allowed to enter. Clay can be seen, and 
renders the water turbid, but organic matter and bacteria may be 
present in water which looks perfectly clear. 

Purification from Suspended Matter. — The suspended 
impurities may be removed by filtration. On a large scale, beds of 
gravel are employed, but this treatment will not remove all bac- 
teria. In many cases small amounts of alum, or alum and lime, 
or ferrous sulphate (copperas) and lime, are added. These pro- 
duce slimy precipitates which assist in the elimination of fine, sus- 
pended inorganic and organic matter, including practically all the 
bacteria. This is called the coagulation treatment (q.v) . The whole 
suspended matter is then allowed to settle, which it does very 
quickly, in large reservoirs. The remaining organisms may be 
destroyed by adding a little bleaching powder (q.v.), before the 
water is distributed. Ozone and ultra-violet light are used for 
the same purpose. 



92 



COLLEGE CHEMISTRY 



In the household, the Pasteur filter is the most compact and 
efficient appliance. The water enters at the top (Fig. 50), and is 
forced inwards by its own pressure through the pores of a cylinder 
of unglazed porcelain. The cylinder must be taken 
out, and its exterior cleaned daily with a brush, to 
remove the mud and organisms which collect on its 
outer surface. If this is not done, the organisms 
multiply and soon the filter pollutes the water instead 
of purifying it. 

Most organisms can be killed by boiling the un- 
filtered water for 10 or 15 minutes, although a second 
boiling is needed in the case of some. 

Purification from Dissolved Matter. — Filtra- 
tion does not remove dissolved matter, and therefore 
does not soften hard water (q.v.). 

Pure water for chemical purposes is prepared by 
distillation and, in fact, liquids other than water are 
usually purified by the same process (Fig. 51). The 
steam is condensed by cold water circulating in the 
jacket, and contains only gases dissolved from the 
air. Dissolved solids remain in the flask. Distilled water quickly 
dissolves traces of glass or porcelain, so that the purest water is 
obtained by using quartz or platinum for the condenser tube and 
receiving vessel. Tin is the best of the less expensive materials. 

Chemical Properties of Water. — Water is so very frequently 
used in chemical experiments in which it is a mere mechanical ad- 
junct, that the beginner has difficulty in distinguishing the cases in 
which it has itself taken part in the chemical interaction. The 
four kinds of chemical activity which it shows should therefore 
receive careful notice: 

1. Water is a relatively stable substance. 

2. It combines with many oxides, forming bases or acids. 

3. It combines with many substances, chiefly salts, forming 
hydrates. 

4. It interacts with some substances in a way described as hy- 
drolysis. This property will not be discussed until a characteristic 
case is encountered. 



. 



Fig. 50. 



WATER 



93 



Water a Stable Compound: [^Dissociation. — In the case of 
a compound, the first chemical property to be given is always, 
whether the substance is relatively stable or unstable. Usually the 




Fia. 51. 

specification is in terms of the temperature required to produce 
noticeable decomposition. Thus, potassium chlorate gives off 
oxygen at a low red heat. Now, water vapor, when heated, is 
progressively decomposed into hydrogen and oxygen, yet at 2000° 
the decomposition reaches only 1.8 per cent, and reunion occurs as 
the temperature is lowered. The two arrows in the equation indi- 
cate that the action may proceed in either direction — is reversible: 

H 2 0<=±2H + 0. 

A decomposition which thus proceeds at higher temperatures, 
while at lower temperatures combination of the constituents can 
take place, is called a dissociation. The decomposition of potas- 
sium chlorate (p. 27) is not a dissociation because it is not revers- 
ible; oxygen gas will not under any known circumstances unite with 
potassium chloride. 



94 COLLEGE CHEMISTRY 

Union of Water with Oxides. — 1. Sodium oxide (Na^O) 
unites violently with water to form sodium hydroxide: 

Na 2 + H 2 0-+2NaOH. 

The slaking of quicklime is a more familiar action of the same kind: 
CaO + H 2 0->Ca(OH) 2 . 

No other products are formed. The clouds of condensing steam 
produced in the second instance are due to evaporation of a part 
of the water by the heat produced in the formation of calcium 
hydroxide. The aqueous solutions of these two products have a 
soapy feeling, and turn red litmus (a vegetable extract) blue, and 
the substances therefore belong to the class of alkalies or bases. 
Very many hydroxides, which are of the same nature, for example 
ferric hydroxide Fe(OH) 3 and tin hydroxide Sn(OH) 2 , are formed 
so slowly by direct union of the oxide and water that they are 
always prepared in other ways. The oxides which, with water, 
form bases are called basic oxides. 

2. Some oxides, although they unite with water, give acids, 
which are products of an entirely different character. Phosphorus 
pentoxide (p. 32) and sulphur dioxide are of this class and yield 
phosphoric acid and sulphurous acid. Such oxides are commonly 
called the anhydrides (Gk., without water) of their respective acids. 
They are called also acidic oxides: 

P 2 5 + 3H 2 0->2H 3 P0 4 . 
S0 2 + H 2 0-*H 2 S0 3 . 

The acids are sour in taste and turn blue litmus red. 

These two classes of final products are so different that we make 
the distinction the basis for classification of the elements present 
in the original oxides. The elements, like sodium and iron, whose 
oxides give bases, are called metallic elements; those, like phos- 
phorus, whose oxides give acids, are called non-metallic elements. 
The distinguishing words are selected because the division corre- 
sponds, in a general way at least, with the separation into two sets 
to which merely physical examination of the elementary substances 
would lead. 



WATER 



95 




Fig. 52. 



Hydrates. — Many substances when dissolved in water and 
recovered by spontaneous evaporation of the solvent enter into 
combination with the liquid. The products, which are solids, are 
called hydrates. That they are regular chemical compounds is 
shown by the following two facts: (1) These compounds show 
definite chemical composition expressible by formulae in terms of 
chemical unit weights (atomic weights) of the constituent ele- 
ments. The proportions in solutions and other physical aggre- 
gations, except by chance, cannot be expressed by means of 
formulae. (2) A hydrate has physical properties entirely different 
from those of the water (or ice) and the 
other substance used in preparing it. It is 
a typical compound, formed by the first 
variety of chemical change (p. 7). Thus, 
cupric sulphate, often called anhydrous 
cupric sulphate to distinguish it from the 
compound with water, is a white substance 
crystallizing in shining, colorless, needle- 
like prisms. The pentahydrate (blue-stone 
or blue vitriol) which crystallizes from the 
aqueous solution, is blue in color, and forms larger but much less 
symmetrical (asymmetric or triclinic) crystals (Fig. 52) : 

CuS0 4 + 5H 2 <=± CuS0 4 ,5H 2 0. 

The chemical properties show hydrates to be relatively unstable. 
When heated, the hydrates, as a rule, lose none of the constituents 
of the original compound, but only the water, in the form of vapor. 
When melted, or when dissolved in water, the hydrates are disso- 
ciated (p. 93) into water and the original substance. The aqueous 
solutions made from the anhydrous substances and from the hy- 
drates have identical physical and chemical properties. Hence 
the cheaper of the two forms is generally purchased, and many 
of the chemicals used in the laboratory are in the form of 
hydrates. 

In consequence of the ease with which hydrates give up water 
we write their formulae (e.g., CuS0 4 ,5H 2 0) so that the water and 
original substance are separate. A formula thus modified, so as 
to show some favorite mode of behavior of the substance, is called 
a reaction formula. The formula Hi CuSO 9 , which would show the 



96 COLLEGE CHEMISTRY 

same proportions by weight, is never employed, because its use 
would disguise the relation of the substance to cupric sulphate. 

The Dissociation of Hydrates, Efflorescence, — The less 
stable hydrates dissociate very readily. Thus the decahydrate 
of sodium sulphate, Na 2 SO 4 ,10H 2 O (Glauber's salt), loses all the 
water it contains (effloresces) when simply kept in an open vessel. 
When kept in a closed bottle, a very little of it loses water, and 
then the decomposition ceases. The cause of this we discover 
when a crystal of the hydrate is placed above mercury, like the ice 
or water in Fig. 49 (p. 87). It shows an aqueous tension which we 
can measure. At 9° the value of this is 5.5 mm. As its tempera- 
ture is raised, the tension increases. When the temperature is 
lowered, on the other hand, the tension diminishes, the mercury 
rises, and a part of the water enters into combination again. 
Different hydrates show different aqueous tensions at the same 
temperature. For example, at 30°, that of water itself is 31.5 mm., 
strontium chloride SrCl 2 ,6H 2 0, 11.5 mm.; cupric sulphate CuSC>4, 
5H 2 0, 12.5 mm.; barium chloride BaCl 2 ,2H 2 0, 4 mm. 

In view of these facts, we perceive that loss of water by efflores- 
cence is like evaporation, excepting that it is a chemical decompo- 
sition and not a physical process. Those hydrates which, like 
Glauber's salt and washing soda Na 2 CO3,10H 2 O, have a vapor ten- 
sion approaching that of water itself, lose their water at ordinary 
temperatures at a rapid pace. Now, atmospheric air is usually 
less than two-thirds saturated with water vapor, and the partial 
pressure (p. 72) of this vapor opposes the dissociation and tends to 
prevent the liberation of the water. Thus at 9°, the vapor tension 
of water being 8.6 mm., the average vapor pressure of water in 
the atmosphere will be about 5 mm. Any hydrate with a greater 
aqueous tension than 5 mm. at 9°, such as Glauber's salt, will 
therefore decompose spontaneously in an open vessel. But those 
with a lower vapor tension, such as the pentahydrate of cupric 
sulphate with a tension of 2 mm. at 9°, will not do so. Granular 
calcium chloride CaCl 2 ,2H 2 is used in drying gases because it 
has an exceedingly low tension of water vapor, and combines with 
water vapor to form CaCl 2 ,6H 2 0. 

The water of hydration is known colloquially in chemistry as 
water of crystallization. The term was introduced when it was first 



WATER 97 

observed that a hydrate, in decomposing, crumbles and loses its 
original crystalline form. But the phrase is misleading. Sulphur, 
potassium chlorate, and thousands of other substances are crys- 
talline, yet do not contain the elements of water. All pure chemi- 
cal substances, in solid form, when in stable physical condition, are 
crystalline. Amorphous (i.e., non-crystalline) substances, like wax 
and glass, are supercooled liquids. 

How Formulae and Equations are Obtained. — In the last 
few pages several formulae (e.g., of hydrates) and several new equa- 
tions have been given. How do we know what to set down in 
making an equation? We cannot learn this by simply writing 
formulae on a piece of paper. In each case, experiments must be 
made in the laboratory. For example, how do we know that the 
common hydrate of cupric sulphate has the formula CuS0 4 ,5H 2 0, 
and not CuS0 4 ,H 2 0? We must make a quantitative experiment. 
We weigh a porcelain dish or crucible, first empty, and then with 
a little of the hydrate. Suppose the difference in weight to be 2.05 
g. (= weight of hydrate). We then heat the dish and contents, 
until the water is driven out, and weigh again. The difference 
is now only 1.31 g. (wt. of anhydrous cupric sulphate). The 
water, therefore, weighed 2.05 — 1.31 = 0.74 g. Assuming that 
we know the formulae (compositions) of cupric sulphate and of 
water, we obtain their formula-weights: CuS0 4 = 63.57 + 32.06 
+ 4 X 16 = 159.63; andH 2 = 2 X 1.008 + 16 = 18.016. The 
formula must be CuS0 4 ,zH 2 0. Also 

159.63 :x X 18.016 :: 1.31 : 0.74. 

Solving for z, we have z X 18.016 X 1.31 = 159.63 X 0.74, or 
x = 159.63 X 0.74/18.016 X 1.31 = 5.00. The formula is there- 
fore CuS0 4 ,5H 2 0, and the equation for the decomposition: 

CuS0 4 ,5H 2 -* CuS0 4 + 5H 2 0. 

To make an equation, we must note what substances are taken, 
and recognize by their properties all the substances produced. If 
all the substances are well known, and we can find their formulae 
in a book, we can make the equation at once. If we cannot find 
the formulae, we make measurements to determine the proportions 
by weight, calculate the formulae, and then make the equation. 



98 



COLLEGE CHEMISTRY 



Composition of Water. — The proportion of hydrogen to oxy- 
gen, in water, by weight, is 2 : 15.879, or 2.016 : 16. The propor- 
tion by volume is 2.0027 volumes of hydrogen to 1 volume of 
oxygen. That the proportion by volume is very close to 2 : 1 may 
easily be shown by mixing hydrogen and oxygen in this propor- 
tion, in a strong tube, and exploding the mixture by means of a 
spark from an induction coil. The resulting steam condenses and 
the whole gas vanishes. If different proportions are used, the 
excess of one of the gases remains uncombined. 



Gay-Lussac's Law of Combining Volumes. 

— The almost mathematical exactness with which 
small integers express this proportion is not a mere 
coincidence. Whenever gases unite, or gaseous prod- 
ucts are formed, the proportions by volume (meas- 
ured at the same temperature and pressure) of all 
the gaseous bodies concerned can be represented 
very accurately by ratios of small integers. This is 
called Gay-Lussac's law of combining volumes (1808). 
Thus, when the above experiment is carried out 
at 100°, in order that the product, water, may be 
gaseous also, it is found that the three volumes of 
the constituents give almost exactly two volumes 
of steam. For example, 15 c.c. of hydrogen and 
7.5 c.c. of oxygen give 15 c.c. of steam. Of course 
the hydrogen, oxygen, and steam must be measured 
at the same pressure, and the temperature must 
remain constant (100°) during the experiment. 
Proper manipulation secures the former, and a 
jacket filled with steam (Fig. 53) the latter con- 
dition. Strips of paper, 1, 2, and 3, are pasted on 
the jacket in such a way that equal lengths of 
the eudiometer, in this case a straight one, are 
laid off. The three divisions being filled with a mixture of hydro- 
gen and oxygen in the proper proportions, the gas, after the 
explosion, shrinks so as to occupy, at the same pressure, only two 
of them. 



Fig. 53. 



WATER 99 

Exercises. — 1. Name some familiar transitions (p. 86) from 
one physical state to another. 

2. What evidence is there in the common behavior of ether 
and chloroform to show that these liquids have high vapor tensions? 

3. If the pressure of the steam in a boiler is ten atmospheres, 
at what temperature is the water boiling (p. 88)? 

4. How many grams of water could be heated from 20° to 100° 
by the heat required to melt 1 kg. of ice at 0°? 

5. What do you infer from the fact that alum and washing soda 
lose their water of hydration when left in open vessels, while 
gypsum does not? 

6. Which fact shows most conclusively that hydrates are true 
chemical compounds? 

7. Gypsum is a hydrate of calcium sulphate (CaSO^. If 6 g. 
of gypsum, when heated, lose 1.256 g. of water, what is the formula 
of the hydrate? 

8. In what ways does a hydrate differ from, (a) a solution, (b) an 
hydroxide? 

9. Should you expect to find any difference, in respect to chemi- 
cal activity, between the three forms of water? Have we had any 
experimental confirmation, or the reverse, of this conclusion (p. 
51)? 

10. Name some crystalline substances which are not used, or 
do not occur in the form of hydrates. 

11. Define the purposes for which evaporation and distillation 
are used. 



CHAPTER IX 

MOLECULAR WEIGHTS AND ATOMIC WEIGHTS 

Gay-Lussac's law (p. 98) shows that, when substances are 
measured in the gaseous condition, and by volume (not by weight), 
the proportions in which they combine can be represented by small 
whole numbers, such as 2 : 1, or 1 : 1, or 2 : 3. The numbers are 
much simpler than when proportions by weight are employed. 
Thus, lead and oxygen combined in the proportions 100 : 7.72 by 
weight. It would seem, therefore, that the shortest route to 
simple methods for expressing combining proportions must lie 
through a study of volumes of gases and vapors. 

Molecular Weights 

The Chemical Unit of Volume for Gases: 22 A Liters, — 

The first thing we require is a suitable unit volume. In making a 
choice, we have to keep in mind the fact that many substances 
cannot easily be converted into vapor, and that therefore meas- 
urement of gaseous volumes cannot entirely displace the measure- 
ment of weights. The measurement of gaseous volumes is only 
to furnish the key to the system. Hence, in choosing our unit 
volume of gas, we must choose one which bears a simple relation 
to our units of weight. Now, the unit of volume chosen is that of 
32 grams of oxygen, which, at 0° and 760 mm. pressure, is 22.4 liters. 
At this stage, it may appear that this is an unduly large unit 

— that 16 grams of oxygen, occupying 11.2 liters, might have 
sufficed. As we proceed, however, we shall find that a smaller 
unit than 22.4 liters leads to a number less than unity for the 
atomic weight of hydrogen. There is no theoretical or chemical 
objection to a unit involving an atomic weight for hydrogen that 
is less than 1, but chemists are unanimous in preferring to have an 
arithmetical unit in the scale, simply as a matter of convenience. 
So, reasoning back from this decision, they have found it necessary 

— we shall perceive the reason presently — to choose 22.4 liters in 
the gaseous condition as the unit quantity of a substance. 

100 



MOLECULAR WEIGHTS 



101 



We shall understand what follows much more readily if we have 
before us, in our minds at least or, better still, in the form of a 
wooden box, a representation of this unit volume (Fig. 54). A 
cube 11.1 inches in height holds 22.4 liters. It is to be under- 
stood that under conditions other than 
0° and 760 mm., this unit volume 
changes in accordance with the laws 
of gases. In this way, it always con- 
tains the same quantity of a given kind 
of gas. In what follows, the standard 
conditions are assumed, unless other 
conditions are specifically mentioned. 



G.M.V. 

22.4 LITERS 



Fig. 54. 

Weights Occupying the Unit Volume — 22.4 Liters. — 

In order that we may keep in touch with the weights, the following 
table gives the weights of equal volumes of several gases and 
vapors. The first column contains the weights of 1 liter. The 
one-thousandth part of each of these weights is the density * 
(p. 73) of the gas (weight of 1 c.c). The experimental method 
of measuring the weight of 1 liter of a gas has already been de- 
scribed (p. 73). In the second column are the weights of 22.4 
liters, obtained by multiplying the values in the first column by 
22.4. It will be observed that the weights of equal volumes of 
the gases cover a wide range of values from 2 for hydrogen (col. 3) 
to 271.5 for mercuric chloride. 



Gases or Vapors. 



Hydrogen . . . . 

Oxygen 

Chlorine . . . . 
Hydrogen chloride 
Carbon dioxide . , 

Water 

Mercury . . . . 
Mercuric chloride . 
Air 



Weight* of One 


Weight of 22.4 


Liter, 0° and 760 


Liters (Molecular 


mm. 


Weight). 


0.090 


2.016 


1.429 


32.00 


3.166 


70.92 


1.628 


36.468 


1.965 


44.00 


0.8045 


18.016 


8.932 


200.6 


12.097 


271.52 


1.293 


28.955 



* These weights are corrected for the two deviations from the laws of gases (pp. 78-79). 

* Sometimes density is expressed on the basis air = 1. One liter of air 
weighs 1.293 g. Hence, if one liter of a gas weighs 3.6 g., its density, air = 1, 
is found by the proportion: 1.293 : 1 :: 3.6 : x. 



102 COLLEGE CHEMISTRY 

The values for the vapors of water (b.-p. 100°), mercury (b.-p. 
357°), and mercuric chloride (b.-p. 300°) are measured at high 
temperatures and reduced by rule (pp. 71-72) to 0° and 760 mm. 

Molecular Weights. — In this discussion of volumes and of 
weights, we must not overlook the interpretation of our results in 
terms of molecules. The masses of gas we handle are aggregates 
of molecules, and the molecules are physically the real units of 
matter. 

Now, according to Avogadro's law, equal volumes of gases (at 
the same t. and p.) contain equal numbers of molecules. The 
weights in each column of the table are therefore weights of equal 
numbers of molecules. The chemical units in the last column 
show, therefore, the relative weights of the individual molecules of 
the substances named. On this account they are called the mo- 
lecular weights of the respective substances. 

Since the 22.4-liter volume holds 32 grams of oxygen and 2.016 
grams of hydrogen — the gram being used throughout — this 
volume is called the gram-molecular volume (G.M.V.) and the 
weights just mentioned are the gram-molecular weights. Fre- 
quently, these ponderous terms are shortened to molar volume and 
molar weight, and the latter even to mole. Thus, a mole of chlo- 
rine is 70,92 g. of the simple substance and a mole of hydrogen 
chloride is 36.468 g. of the compound.* 

Measurement of Molar Weights (Moles). — We may now 
state, in brief, the method of finding the molar (gram-molecular) 
weight of a substance thus: Weigh a known volume of the substance, 

* A common question is: Do not molecules of different substances differ 
in size, and will not the numbers required to fill the G.M.V. therefore be differ- 
ent? The answer is that the molecules are all so small compared with the 
spaces between them (at 760 mm.) that the distances from surface to surface 
are practically the same as from center to center. A G.M.V. of oxygen, 
when liquefied, gives less than 32 c.c. of liquid oxygen, or less than 1/700 of 
the volume as gas. It is only when gases are so severely compressed that the 
nearness of the molecules to one another approaches that found in the liquid 
condition that the effects of the bulk of the molecules become conspicuous, 
and a difference in the behavior of different gases is noticeable. But in the 
work discussed in this chapter, pressures over one atmosphere are intention- 
ally avoided. 



ATOMIC WEIGHTS 103 

at any temperature and pressure at which it is gaseous, reduce this 
volume by rule to 0° and 760 mm., and calculate by proportion the 
weight of 22.4 liters (see Exercises 1, 2, 3, 5). 

That quantity of each gaseous substance which at 0° and 760 mm. 
would fill the G.M.V. cube is the unit quantity of the substance for 
all theoretical purposes in chemistry. It represents the relative 
weight of the molecules of the substance. We shall employ it 
presently for the purpose of determining the relative weights of 
atoms, or atomic weights. 

The Number of Molecules in a Mole. — The molecular 
weight or mole of a substance is not the weight of a single molecule. 
It is only the relative weight of the molecule of the substance. It 
is, however, the weight in grams of a fixed number of molecules, 
for 22.4 liters (or any other volume) contains equal numbers of 
molecules of different gases. The actual number has been deter- 
mined. Thus Jean Perrin found values by several experimental 
methods which ranged between 5.9 X 10 23 (that is, 59 followed by 
22 ciphers) and 6.9 X 10 23 . Rutherford, using an entirely differ- 
ent plan, obtained 5.7 X 10 23 for the gas helium. The value 
which is accepted as most accurate was that obtained by R. A. 
Millikan of the University of Chicago, by the use of a still differ- 
ent method, namely, 6.07 X 10 23 (or 6070 2 i). 

Atomic Weights 

Chemical Unit Quantities of Elements. — We are now 

approaching the question of units, in which to express combining 
proportions, from a different view point from that employed in the 
earlier chapters. We were then assigning numbers for the quan- 
tities of the constituent elements of a compound (such as iron : 
oxygen :: 111.68 : 48, p. 9) without any consideration of the 
magnitude of the total weight of the constituents. At that time, 
we had no reason before us to indicate that this total might require 
consideration. We now start by determining and assigning the 
total weight of the compound, and it is our next task to consider 
the subdivision of this total amongst the constituents. Evidently, 
if the unit quantity of the compound has been properly chosen, 
it must be subdivisible into one or more unit quantities, of suit- 
able dimensions, of each element in the compound. Let us now 



104 COLLEGE CHEMISTRY 

set down, and examine the results of such a subdivision in the 
case of several compounds. To be more precise, we take 22.4 
liters of every substance — one cubeful in the gaseous condition 
— as the total quantity. We make an analysis of a sample of 
the substance, in case it is a compound, to ascertain the propor- 
tions in which the elements are present in it. We then divide 
the weight of 22.4 liters of the compound between the different 
elements in the proportion shown by the analysis. 

For example, the cube holds 36.468 g. of hydrogen chloride gas. 
This amount, when decomposed, yields 1.008 g.* of hydrogen and 
35.46 g. of chlorine. 

Another example: Suppose the substance is a liquid, like phos- 
phorus oxychloride. We determine the weight of a measured 
volume of its vapor, at a properly chosen temperature and pres- 
sure, and the result gives us, by calculation, the weight of 22.4 
liters, the molecular weight, viz., 153.38. That is, 153.38 g. of 
the substance would fill the cube, if it could be kept as vapor at 
0° and 760 mm. The analysis shows that this amount of the 
substance contains 31 g. of the element phosphorus, 16 g. of the 
element oxygen, and 106.38 g. of the element chlorine. 

In the following table a few sample results of the process just 
outlined are given. The first column contains the molar weight, 
i.e., the weight of the substance which occupies the G.M.V. cube. 
In the other columns are entered the weights of the various ele- 
ments which together make up the total molar weight. - To sim- 
plify the numbers, the values used are hydrogen 1, phosphorus 
31, mercury 200, instead of 1.008, 31.04, and 200.6, respectively. 

Atomic Weights. — To contain similar data for all the volatile 
compounds of every known element, a huge table, of which this 
might be a small corner, would be required. With such a table at 
hand the atomic weight of each element could promptly be picked 
out. Thus, in the carbon column it would be found that all the 
weights of carbon were either 12 or integral multiples of 12, and 

* It will be observed that if the unit for molecular weights had been less 
than the number of molecules in 22.4 liters of oxygen, then an equal number of 
molecules of hydrogen chloride would have contained less than 1.008 g. of 
hydrogen, and the atomic weight of this element would then have been lesa 
than unity. 



ATOMIC WEIGHTS 



105 





Molar 
Weight. 


Weights of Constituents in Molar Weight. 


Substance. 


® 

I 

>> 

w 


i 

1 

o 


1 

i 

o 


3 

.a 
a 

8 

P4 


a 
o 


>> 

% 
p 

B 

5 


Molecular 
Formula. 


Hydrogen chloride. . . . 
Chlorine dioxide 


36.46 

67.46 
137.38 
153.38 
284 

34 * 

18 

16 

26 

28 

30 

60 
235.46 
270.92 


1 

*3" 

2 
4 
2 
4 
2 
4 


35.46 

35.46 

106.38 

106.38 

35^46 
70.92 










HC1 


32 

i6' 

160 

i6* 
i6 

32 








C10 2 


Phosphorus trichloride. . 


31 
31 

124 
31 






PC1 3 


Phosphorus oxychloride . 
Phosphorus pentoxide . . 

Phosphine 

Water 






POCl 8 






P4O10 






PH3 






H 2 


Methane 

Acetylene 

Ethylene 

Formaldehyde 

Acetic acid 

Mercurous chloride . . . 




12 

24 
24 
12 
24 


200 ' 
200 


CH4 

C2H2 
C2H4 

CH 2 o 

C2-H402 

HgCl 
HgCl 2 


Mercuric chloride .... 

















this is therefore the most convenient unit weight (and therefore 
the atomic weight) of carbon. Similarly, the atomic weight of 
oxygen is 16,* of phosphorus 31, of mercury 200 (see Exercise 4). 

The fact that all the numbers in any one column turn out to be 
even multiples of a single number need not seem mysterious. The 
molecule of every compound containing chlorine must contain 
one, two, three, or some other whole number of chlorine atoms, for 
chlorine atoms, like other atoms, do not furnish fractions of atoms 
in any cases of combination. Now, the weight of chlorine in 
6O7O21 atoms, assuming one atom of chlorine to each molecule in 
22.4 liters of some gas containing chlorine, must be 35.46 g. 
Hence, if the weight of chlorine in 22.4 liters (6070 2 i molecules) 
of the compound differs from 35.46 g., it can do so only because 
there are two atoms of chlorine per molecule, giving 2 X 35.46 g., 
or three atoms giving 3 X 35.46 g. of chlorine, and so forth. Thus 
the quantities of chlorine in the G.M.V. of all compounds of 
chlorine must be a multiple of 35.46 by unity or some other integer. 

When the atomic weights have finally been selected, we can go 
through the table and change all the numbers into multiples of the 

* The difference between the unit quantity of oxygen in compounds 
(namely 16) and the unit quantity of free oxygen (32) will be discussed presently. 



106 COLLEGE CHEMISTET 

chosen atomic weights. Thus, for 70.92 we write 2 X 35.46, and 
for 106.38 we write 3 X 35.46, and so forth. The reader should 
prepare such a modification of the table. With this new form of 
the table before us, we can, finally, replace the atomic weights by 
the symbols which stand for them, writing, for 35.46, CI, for 
2 X 35.46, CI2, and so forth. The results of doing this in each line, 
i.e., for each substance, are collected at the ends of the lines in the 
last column of the table. The reader should himself repeat the 
substitutions of the symbols, and so verify the formulae given. 
These formulae, since they are based on the molecular weights, 
in such a way that when the numerical values are substituted for 
the symbols the total restores to us the molecular weight, are 
called molecular formulae. 

As a definition, the atomic weight of an element may be stated 
to be: The smallest of the weights of the element found in the 
molecular weights of all its volatile compounds, so far as these have 
been examined. 

It is hardly necessary to add that the atomic weights, found as 
described above, are equally serviceable in expressing the compo- 
sitions of compounds which are not volatile. The atoms in non- 
volatile compounds are identical in properties with the atoms of 
the same elements in volatile compounds. If an element gives no 
volatile compounds, other methods of fixing its atomic weight are 
available (see Dulong and Petit's law, p. 108). 

Although in this section, as well as elsewhere, we have empha- 
sized the fact that atoms are not divided into parts, this must not 
be taken to mean that atoms are incapable of being broken up. 
It means only that in ordinary chemical changes, the atoms com- 
bine and separate as wholes. Indeed, we now know that the atom 
of radium (q.v.) gives off atoms of helium, and leaves an atom of 
lead, and that the atoms of one or two other elements disintegrate 
in a similar way. Some day means of breaking up any or all 
kinds of atoms may be discovered. 

Many chemists have contributed to the determination and re- 
vision of the atomic weights. The Swedish chemist, Berzelius, 
devoted many years to the accurate measurement of combining 
proportions. Stas, a Belgian (1860-1870), made a number of 
determinations with great exactness. Morley's (1895) value 
for combining proportions of hydrogen and oxygen alone repre- 



ATOMIC WEIGHTS 107 

sented several years of work. T. W. Richards of Harvard Uni- 
versity has recently carried many of the values to a higher degree 
of accuracy. 

Why 22.4 Liters was Chosen as the Unit Volume. — We 

can now see why the volume occupied by 32 g. of oxygen, namely, 
22.4 liters, was taken as the standard for the scale of molecular 
quantities. This gave us, for example, 36.468 g. as the weight of 
22.4 1. of hydrogen chloride, which in turn contains 1.008 g. of 
hydrogen. A smaller weight of oxygen, with correspondingly 
smaller standard volume, would have held an amount of hydrogen 
chloride (and of other compounds containing one atom of hydro- 
gen per molecule) which would have been less than 1 gram. The 
choice was made to secure something close to an arithmetical 
unit in the scale. 

Advantages of Atomic Weights. — Although the method 
of selecting atomic weights involves rather complex reasoning, 
these weights repay the trouble, because they represent the rela- 
tive weights of the atoms themselves. They are thus much more 
valuable in helping us to understand chemical behavior and in 
enabling us to classify the phenomena of chemistry than would be 
any other units of weight we might have chosen. The following 
are some of the advantages they offer: 

1. The atomic weight of an element has but one value, and 
this value is definitely determinable. The advantages of using 
Avogadro's principle (1811), and taking a unit volume of gas as 
the basis of chemical units, were not perceived by chemists until 
Cannizzaro, in 1858, succeeded in setting them forth in a con- 
vincing manner. Previous to that time, different chemists used 
different unit weights for the same element, and therefore assigned 
different formulae to the same compound, and much confusion ex- 
isted. After 1858 chemists united upon the present values for 
atomic weights. 

2. The atomic weight of an element has a valence (p. 61), while 
equivalents are equi-valent. While valence is a helpful conception 
in all branches of chemistry, organic chemistry is especially in- 
debted to the conception of the quadrivalence of carbon for much 
of its development and most of its organization. The full illus- 
tration of this point is beyond the limits of the present book. 



108 



COLLEGE CHEMISTRY 



3. The periodic system (q.v.), the basis of a plan for classifying 
the properties of all chemical substances, is founded upon the 
atomic weights. 

4. Dulong and Petit's law is based upon atomic weights. This 
law furnishes also an alternative means of determining atomic 
weights that has frequently rendered valuable service, and on this 
account forms the subject of the next section. 

Dulong and Petit'' s Law, an Alternative Means of Deter- 
mining Atomic Weights. — It was first pointed out (1818) by 
Dulong and Petit, of the Ecole Poly technique in Paris, that when 
the atomic weights of the elements were multiplied by the specific 
heats of the simple substances in the solid condition, the products 
were approximately the same in all cases. In other words, the spe- 
cific heats are inversely proportional to the magnitudes of the 
atomic weights. The table, in which round numbers have been 
used for the atomic weights, shows that the product lies usually 
between 6 and 7, averaging about 6.4: 



Element. 


Atomic 
Wt. 


Sp. Ht. 


Prod- 
uct. 


Element. 


Atomic 

Wt. 


Sp. Ht. 


Prod- 
uct. 


Lithium . . . 
Sodium . . . 
Magnesium . 
Silicon . . . 
Phosphorus 

(Yellow) . 
Calcium. . . 


7 

23 

24.3 

28.3 

31 
40 


0.94 
0.29 
0.245 
0.16 

0.19 
0.170 


6.6 

6.7 
6.0 
4.5 

5.9 

6.8 


Iron 

Zinc 

Bromine (Solid) 

Gold 

Mercury (Solid) 
Uranium . . . 


56 

65.4 

80 
197 
200 
238.5 


0.112 

0.093 

0.084 

0.032 

0.0335 

0.0276 


6.3 
6.1 

6.7 
6.3 
6.7 
6.6 



Another way of expressing this law will give it greater chemical 
significance. The specific heats are the amounts of heat required 
to raise one gram, that is one physical unit, of each element through 
one degree. When we multiply this by the atomic weight, we 
obtain the amount of heat required to raise one gram-atomic weight 
of the element, that is, one chemical unit, through one degree. 
The values of this product are approximately equal. Since there 
are equal numbers of atoms in one gram-atomic weight of each ele- 
ment, it follows that: Equal amounts of heat raise equal numbers 
of atoms of all elements in the solid form through equal intervals of 
temperature. 



MOLECULAR EQUATIONS 109 

It will be seen at once that, although the law of Dulong and 
Petit is purely empirical, it may nevertheless be used for fixing the 
atomic weight of an element of which no volatile compounds are 
known. We can always measure that weight of such an element 
which combines with one atomic weight of another element. Since 
the elements concerned must combine atom for atom, or in some 
simple ratio such as 1 : 2 or 2 : 3, it follows that the weight found 
is either the atomic weight or some multiple or submultiple of it 
by a whole number. When, therefore, we multiply this weight 
by the specific heat, we discover at once whether the product is 
6.4 or some simple fraction or multiple of this number. For ex- 
ample, suppose the atomic weight of calcium to be unknown. 
We find by analysis that calcium chloride contains 20 parts of 
calcium combined with 35.46 parts (the atomic weight) of chlo- 
rine. Now the specific heat of solid, metallic calcium is 0.170. 
This number multiplied by 20 gives as the product 3.4. Evi- 
dently, therefore, the atomic weight is not 20, but 40, for the 
product, 6.8, then agrees fairly well with the average for other 
elements. 

Molecular Formulae 

Molecular Formulae of Compounds. — If the molar formulae 
in the table (p. 105) be examined it will be observed that several 
are not in their simplest terms. Thus, the formula of acetylene is 
C2H2. The formula CH would represent the composition of the 
substance equally well, for 12 : 1 is the same as 24 : 2. But the 
formula CH gives a total of only 13, while C 2 H 2 shows the total 
weight of the molecule to be 26 and records for us therefore the 
weight of the G.M.V., as well as the composition of the substance. 
We shall find this additional property, peculiar to the molecular 
formula, to be a feature of the greatest practical value. Some 
of the practical uses of this improvement in our formulae will be 
illustrated in this chapter, and there is an example of one of them 
in the table itself. Thus, the molecular formula of acetic acid is 
C2H4O2, and not the simpler, identical proportion CH 2 0. The 
latter is the molecular formula of a totally different substance, 
formaldehyde, now much used as a disinfectant. The vapor of 
this substance has only half the density of acetic acid vapor, and 
this fact, recorded in the formula, helps to remind us that the 



110 COLLEGE CHEMISTEY 

substances are different. Still another substance of the same 
composition is grape sugar (dextrose), CeH^Oe- In addition to 
this and other practical advantages, molecular formulae satisfy also 
the claim of logical consistency. If the symbols represent the 
atomic weights, the formulae should be constructed so as to rep- 
resent the molecular weights. 

Molecular formulae like C 2 H 2 and C2H4O2 are easily interpreted 
in terms of the atomic hypothesis. C represents one atom of 
carbon and H one atom of hydrogen. But there is no reason 
why a molecule of acetylene should not contain two atoms of each 
kind. Similarly, the molecule of formaldehyde contains four atoms 
(CH 2 0), and one of acetic acid eight atoms (C2H4O2), and one of 
dextrose twenty-four atoms (C6H12O6), although the relative num- 
bers of each kind are the same. Indeed this hypothesis helps to 
clear the matter up, for chemists go so far as to account for the 
chemical behavior of the substances by an imagined geometrical 
arrangement of the atoms in their molecules, and these three 
kinds of molecules are supposed to differ in structure as well as in 
the number of atoms they contain. 

The Molecular Weights and Formulae of Elementary 
Substances. — The following table gives the densities of some 
elementary substances, including those of which the substances 
previously discussed are compounds. The first column shows the 
atomic weight, which in each case is the minimum weight of the 
element found in a G.M.V. of any compound. For example, 16 g. 
of oxygen and 35.46 g. of chlorine are the weights in the amounts 
of water vapor and hydrogen chloride, respectively, which fill the 
cube (22.4 liters). The symbol, in the next column, stands for 
this quantity and occurs in many formulae, such as H 2 and HO. 
It represents the combining unit or atom. In the third column is 
given the weight of the free, elementary substance which fills the 
G.M.V. and is the molecular weight. It shows the weight of the 
molecule relative to the weights of the other molecules in the same 
column, and to the weights of the atoms in the first column. In 
the last two columns are given the molecular weights resolved 
into multiples of the atomic weights and the corresponding 
formulae. 



MOLECULAK EQUATIONS 



111 



Element. 



Oxygen . . 
Hydrogen . 
Chlorine . 
Phosphorus 
Mercury . 
Ozone . . 
Cadmium . 
Potassium. 
Sodium . . 
Zinc . . . 



Atomic 


Sym- 


Weight in 


Weight. 


bol. 


G.M.V. 


16.00 


O 


32.00 


1.008 


H 


2.016 


35.46 


CI 


70.92 


31.04 


P 


124.16 


200.6 


Hg 


200.6 


16.00 


O 


48.00 


112.4 


Cd 


112.4 


39.10 


K 


39.10 


23.00 


Na 


23.00 


65.37 


Zn 


65.37 



Weight in 

G.M.V. Fac- 

torized. 



2X16.00 
2X1.008 
2x35.46 
4X31.04 
1X200.6 
3X16.00 
1X112.4 
1X39.10 
1X23.00 
1X65.37 



Formula 

of Free 

Element. 



o 2 

H 2 

Cl» 

P 4 

Hg 

3 

Cd 

K 

Na 

Zn 



The reader cannot fail to note a striking peculiarity. In the case 
of chlorine the molecular weight is 70.92, while the atomic weight 
is 35.46. With hydrogen and oxygen, also, the molecular weight 
contains two atomic w r eights. Yet this is not a general rule, for 
with mercury and several other elements the molecular and atomic 
weights are alike, while with phosphorus the molecular is four times 
the atomic w-eight. Evidently there is no rule, and each element 
has to be subjected to separate experimental study. The result is 
that for free, elementary chlorine we use the molecular formula Cl 2 , for 
free hydrogen H 2 , for elementary, uncombined oxygen the formula 2 . 
For a substance like phosphorus, which is not a gas and is not often 
used as a vapor, the formula P is commonly employed by chemists, 
to avoid the larger coefficients which P 4 introduces into equations, 
although theoretically the latter formula would be the strictly cor- 
rect one. 

The case of oxygen demonstrates clearly the necessity of using 
molecular formulae, even for simple substances. The table shows 
two substances containing nothing but oxygen. Ozone (q.v.) has 
a molecular weight 48, being a gas exactly one-half heavier than 
ordinary oxygen. Its formula, therefore, is O3, while that of oxy- 
gen is 2 . Oxygen and ozone are entirely different chemical indi- 
viduals. The latter has, for example, a strong odor and is much 
more active. Thus polished silver remains bright indefinitely in 
pure oxygen, but oxidizes quickly when placed in ozone. 

To avoid a common error, the reader should note that to learn the 
atomic weight of an element, we do not measure the molecular 



112 COLLEGE CHEMISTRY 

weight of the simple substance. The molecular weight of the ele- 
mentary substance may be a multiple of the atomic weight, and we 
find out whether it is such a multiple only after the atomic weight 
has been determined. The atomic weight is the unit weight used 
in compounds, and can be ascertained only by a study of com- 
pounds. The molecular weight of the free element gives us only 
a value which we know must be a multiple of the atomic weight, by 
1 or some other integer. Mol. Wt. = At. Wt. X x, where # is 1 or 
some other integer. 

Further Discussion of the Molecular Formulae of Ele- 
mentary Substances. — Some further explanation may be re- 
quired, to the end that the reader may be reconciled to accepting 
the formulae Cl 2 , 2 , and so forth. In the first place, he should 
note how these formulae arose. If we accept Avogadro's law, and 
the inference from it to the effect that the weights of equal vol- 
umes of gases are in the same ratio as the weights of their indi- 
vidual molecules, then we cannot escape the conclusion to which 
measuring the relative densities of free chlorine and hydrogen 
chloride, for example, leads. The ratio of their densities is 70.92 : 
36.46. That is to say, the relative weights of a molecule of chlo- 
rine and a molecule of hydrogen chloride stand in this ratio. The 
molecule of chlorine is nearly twice as heavy as the molecule of the 
compound, and there cannot therefore be a whole molecule of chlorine 
in a molecule of hydrogen chloride. In fact, we perceive at once 
that the molecule of hydrogen chloride must contain only half a 
molecule of chlorine (35.46), together with half a molecule of 
hydrogen (1). In other words, if the molecule of free chlorine 
were to be taken as the atom of the element, then the molecule of 
hydrogen chloride would contain only half an atom of chlorine, 
which would be contrary to our definition to take as atoms quan- 
tities which are not divided. So we choose the other horn of the 
dilemma, and say that the specimen of chlorine in the molecule of 
hydrogen chloride is a whole atom and that therefore the amount 
of chlorine in the molecule of free chlorine is two atoms, and its 
formula CI2. Similarly, the weight of hydrogen in the molecule of 
hydrogen chloride is 1.008, while that of the molecule of hydrogen 
is 2.016, so that there are two atoms in the molecule of free hydro- 
gen and its formula is H 2 . Reasoning in like manner from the 



MOLECULAR EQUATIONS 113 

molecular weights of oxygen (32) and water (18) we reach the 
conclusion that the molecule of oxygen is diatomic (0 2 ) . 

The simple fact that hydrogen and oxygen, when mixed, do not 
combine (p. 59) may assist in reconciling us to the diatomic nature 
of their molecules. Some part of the mixture has to be heated 
strongly to start the interaction. Now the molecular formulae, H 2 
and 2 , suggest that each gas is really in combination already (with 
itself), and they therefore explain to some extent the indifference of 
the gases towards one another. If the molecules were free atoms, 
they could not encounter one another continually as they move 
about, and yet escape combination as we observe that they do. 
We may imagine that the primary effect of heating is to decom- 
pose some of the molecules, and liberate hydrogen and oxygen in 
the atomic condition, and that the combination of these atoms 
starts the explosion of the whole mass. 

In the case of hydrogen, the diatomic nature of the molecules 
has been demonstrated by an entirely different method by Lang- 
muir. It has long been known that the conductivity of hydrogen 
for heat is greater than that of any other elementary gas. Thus, 
a wire raised to a white heat in air by means of an electric current, 
cannot be kept at a red heat, even, by the same current in hydro- 
gen. In other gases, heat from the hot wire is used up in accel- 
erating the motion of the molecules of the gas. Langmuir has 
shown, however, that in hydrogen, additional heat is consumed in 
causing decomposition of many of the diatomic molecules into 
single atoms: 

H 2 ^±2H. 

He has measured the percentage of molecules dissociated (at 760 
mm.), and found that it varies from 0.33 per cent at 2000° to 13 per 
cent at 3000° and 34 per cent at 3500°. When the temperature 
falls, the atoms re-combine to diatomic molecules. It may also 
assist in making the matter clear if we note that the atomic weight 
of an element is the unit quantity of that particular variety of 
matter, when it is in combination. The unit quantity of the same 
variety of matter, when in the free state, as a substance, need not 
be the same. We should not expect it to be smaller, but it might 
easily be twice or more times as large. 



114 



COLLEGE CHEMISTRY 



Applications 

Applications: Interactions Between Gases. — According to 
Avogadro's hypothesis, if we filled a succession of vessels of equal 
dimensions with different gases, and could arrest the motion of 
the particles and observe their disposition, we should find that the 
average distance from particle to particle would be the same in 
all cases. This would be true whether our vessels were filled with 
single gases, with homogeneous mixtures, or with gases in layers. 
Such being the case, if any chemical change is brought about in the 
mass which results in a multiplication of the molecules, it is evi- 
dent that the volume will have to increase in order that the spacing 
may remain the same as before. If any chemical action results in 
a diminution of the number of molecules, then a shrinkage must 
take place in order that the spacing may be preserved as before. 
Thus, in a mixture of hydrogen and oxygen, according to our 
hypothesis, when the interaction occurs, the following change 
takes place between neighboring molecules: 

HH + 00 + HH becomes HOH + HOH. 

Since the oxygen molecules, which form a third of the whole, dis- 
appear into the molecules of hydrogen, the tendency to preserve 
spacing results in a diminution of the volume by one-third (p. 98). 
Thus Gay-Lussac's law would have followed as a natural infer- 
ence from Avogadro's law, if the former, being more obvious, had 
not been discovered first. 

If each of the following squares represents a small volume con- 
taining 1000 molecules of gas, then 2000 molecules of hydrogen 
and 1000 molecules of oxygen give 2000 molecules of water vapor. 
We may note, in passing, that, since each molecule of water must 
contain at least one atom of oxygen, at least 2000 atoms of oxygen 
were required, and must have been furnished by the 1000 mole- 
cules of oxygen. Each of these molecules must therefore have 
split into two atoms. 



Hyd. 
1000 
mols. 


+ 


Hyd. 
1000 
mols. 


i 


Ox. 
1000 
mols. 


- 


Aq. 
1000 
mols. 


Aq. 
1000 
mols. 



This method of looking upon chemical interactions between 
gases gives us the nearest sight which we can have of the behavior 



APPLICATIONS OF MOLECULAR EQUATIONS 115 

of the molecules themselves. We cannot perceive the individual 
molecules, but, in consequence of the spatial arrangement which 
they observe, the change in the whole volume of a large aggre- 
gate of molecules enables us to draw conclusions at once in regard 
to the behavior of the single molecules in detail. 

Applications: Molecular Equations. — To utilize the fore- 
going considerations, chemists always employ in their equations the 
molecular formulas for the gases and the easily vaporized substances 
concerned. Thus far, we have used the equation: 

2H + -> H 2 
Weights: 2 X 1.008 16 18.016 

and the information it contained was exhausted when we had placed 
below the symbols the weights for which they stood. But the 
molecular equation is much more instructive. The following 
shows the interpretations to which the molecular equation is subject: 



36.032) g. 



The weights, although doubled, show the same proportions, so 
that questions of weight are answered as easily as before. These 
weights, however, being molecular weights, or multiples thereof, 
can be translated at once into volumes, and questions about volumes 
can also be answered. Finally, the relative numbers of each kind of 
molecules can be read from this equation, for the coefficients in 
front of the formulae represent these numbers. Where no coef- 
ficient is written, 1 is to be understood.* 

Applications: The Making of Molecular Equations, — 

To make a molecular equation, we first make an equation accord- 
ing to the rules already explained (p. 51). An equation like that 
given for the interaction of potassium on water (p. 50) : K + H 2 
— > KOH + H, is the result. Then we adjust the equation so 

* The application of these properties of molecular equations is illustrated 
in Chap. XI (pp. 149-153). If desired, these applications may be taken up 
after the next section. 



2H 2 + 


2 -> 2H 2 


;hts: 2 X 2.016 g. 


32 g. 2 X 18.016 ( 


-ues: 2X22.4 1. 


22.4 1. 2 X 22.4 1. 


:cules: 2 


1 2 



116 COLLEGE CHEMISTRY 

that molecular formulae are used throughout. The hydrogen must 
appear as H 2 , or a multiple of this, in such equations. Hence 
the whole equation must be multiplied by 2 : 

2K + 2H 2 -> 2KOH + H 2 . 

Again, the equation for the preparation of oxygen from potassium 
chlorate: KC10 3 -» KC1 + 30 (p. 27), becomes: 

2KCIO3 -> 2KC1 + 30 2 . 

Every equation containing an odd number of atoms of a substance 
whose molecules are diatomic must be multiplied by 2. Again, 
mercuric oxide decomposes to give mercury vapor and oxygen 
(p. 14), and the molecules of mercury are monatomic and those of 
oxygen diatomic, so we write: 

2HgO -> 2Hg + 2 . 

Finally, the formulae of substances which are solid or liquid, and 
cannot be easily vaporized, are written in the simplest terms. 
Thus, since substances like the copper in the following equation 
are involatile, the molecular weights of such substances are un- 
known, and their molecular formulae likewise: 2Cu + 2 — > 2CuO. 
Furthermore, in the case of substances which can be volatilized, 
although the molecular weights and molecular formulae may there- 
fore be known, we do not usually employ the molecular formulae if 
the substance is not used in the form of vapor in the laboratory. 
Thus, the molecular formula of phosphorus pentoxide is P4O10 
(p. 105). But we generally make, and use, only the solid form, 
and not the vapor, in actual work. Hence the action with water 
is usually written as we have given it (p. 94), rather than in the 
form: P 4 Oio + 6H 2 -+ 4H 3 P0 4 . 

Molecular equations will be used exclusively hereafter. 

Applications: To Cases of Dissociation. — Several gases 
or vapors yield smaller values for their densities, and therefore 
molecular weights, when the densities are measured at higher tem- 
peratures. This indicates that the molecules have become lighter, 
and can only mean that decomposition has taken place in conse- 
quence of the heating. Behavior of this kind is shown both by 
compounds and by simple substances. 



APPLICATIONS OF MOLECULAR EQUATIONS 117 

For example, phosphorus pentachloride PCI5, although a solid, 
can be converted into vapor without much difficulty. Its molec- 
ular weight, if it underwent no chemical change during the vola- 
tilization, would be 31 + 177.3 = 208.3. The density actually 
observed at 300° and 760 mm. pressure gives by calculation not 
much more than half this value. The direct inference from this 
is that the molecules have only half the (average) weight that we 
expected; or, in other words, are twice as numerous as we expected. 
The explanation is found when we examine the nature of the vapor 
more closely. We find that it is a mixture of phosphorus trichlo- 
ride and free chlorine, resulting from a chemical change according 
to the equation : PC1 5 <=± PCI3 + Cl 2 . The low value of the den- 
sity thus tells us that dissociation has taken place. From the 
value of the density at various temperatures, we may even calcu- 
late the proportion of the whole material which is dissociated. At 
300° it is 97 per cent; at 250°, 80 per cent; and at 200°, 48.5 per 
cent. Thus, when the temperature is lowered, progressive re- 
combination takes place and the proportion dissociated becomes 
less. Finally the vapor condenses and yields the original solid. 

Again, sulphur boils at 445°, but can be vaporized at a tempera- 
ture as low as 193°, under very low pressure. At this temperature 
the density of the vapor gives the molecular weight 256 ,(=8X 
32), and the molecular formula S 8 . That is to say, the G.M.V. 
holds 256 g. of the vapor at 193°. At 800°, however, the density 
is only one-fourth as great, and the G.M.V. holds only 64 g. (S 2 ). 
This means that 256 g. now occupy four times as large a volume 
as before, and the increase is additional to the effect of the mere ther- 
mal expansion, which is allowed for in the calculation and elimi- 
nated. Hence the .molecules have dissociated. At 1700° the 
molecular formula is still S 2 , so that this represents the limit of 
dissociation: Sg ^ 4S 2 . When the vapor is cooled, the density 
increases once more and at 193° recovers completely the greater 
value. Similar observations show that phosphorus vapor at 313° 
is all P 4 , but at 1700° one-half of the molecules are P 2 . Iodine 
vapor, up to 700°, is all I 2 . Beyond this temperature the density 
diminishes, and when 1700° is reached the vapor is all I. Thus 
the molecules are diatomic at low temperatures and monatomic at 
high ones. The densities of oxygen, hydrogen, and chlorine are 
not measurably affected by heating to 1700°, so that their dia- 



118 COLLEGE CHEMISTRY 

tomic molecules exist from temperatures far below 0° up to 
1700°, and are evidently very stable. For observations on hydro- 
gen above 1700°, however, see p. 113. 

Applications: Finding the Atomic Weight of a New Ele- 
ment. — By way of reviewing the principles explained in this 
chapter, let us apply them to the imaginary case of a newly dis- 
covered element. The bromide of the element is found to be 
easy of preparation and to be volatile. The bromide contains 
30 per cent of the element (and therefore 70 per cent of bromine), 
and its vapor density referred to air is 11.8. The analysis can 
always be made much more accurately than the measurement o£ 
vapor density, so that the former number is more trustworthy 
than the latter. 

To find the equivalent of the element, that is, the amount com- 
bined with 79.92 parts (the atomic weight) of bromine, we have the 
proportion 70 : 30 :: 79.92 : x, from which x = 34.3. The atomic 
weight must be this, or some small multiple of it. 

The G.M.V. of air weighs 28.955 g. (p. 101). Hence the same 
volume of the vapor of this bromide, which is 11.8 times as heavy as 
air, will weigh 28.955 X 11.8, or 341.67 g. This is therefore the 
molar weight of the compound. 

Now 30 per cent of this is the new element : 

341.67 X 30 ^ 100 = 102.5. 

Now 34.3 parts of the element combined with 79.92 parts"of bro- 
mine. Evidently the atomic weight of the element is 3 X 34.3 = 
102.9, the difference being due to error in determining the density. 
So long as no other volatile compound is known, we adopt this as 
the atomic weight. The rest of the molar weight (239 parts = 
3 X 79.92) is bromine. Thus the formula of the compound is 
ElBr 3 , and from this we see that the element is trivalent. 

In case no volatile compound of the element can be formed, the 
weight combining with 79.92 parts of bromine is measured as before. 
Then some of the free simple substance is made, say by electrol- 
ysis, and its specific heat is determined. The sp. ht. is about 
0.063. Application of Dulong and Petit's law then gives the 
atomic weight. The product 34.3 X 0.063 is equal to 2.161. 
Hence, the equivalent must be multiplied by 3 to give the atomic 



APPLICATIONS OF MOLECULAR EQUATIONS 119 

weight, for this raises the product to 6.48, which is within 
the limits. Thus the value of the atomic weight is 102.9, as 
before. 

Replies to Questions about Difficulties. — The beginner 
always becomes confused over one or more of the points raised by 
the following questions: 

1. Why was 32 g. of oxygen taken as the standard for molecular 
weights, rather than 16 g.? Read p. 107 and footnote to p. 104. 

2. If 2 is the smallest mass of oxygen, why do we have formulaB 
like H 2 and HCIO? 2 is the smallest mass of free oxygen, but 
in combination half as much occurs in many molecules. Read 
pp. 105, 110, and 111. 

3. Why is not the atomic weight of an element ascertained by 
simply measuring the density of the elementary substance? Read 
pp. Ill, last par., and 117, second par. 

4. Can we not deduce the valence of an element from knowing 
the number of atoms in its molecules, and vice versa? Some molec- 
ular formulae and valences are : H 2 X , 2 n , CV, Zn n , also Hg (uni- 
valent and bivalent), P 4 (trivalent and quinquivalent), and Sg 
(bivalent and sexivalent) . There is no relation, either observable 
or to be expected. 

5. Do the molecular weights, oxygen = 32 and hydrogen = 2, 
mean that the molecules of oxygen are larger than are those of 
hydrogen? This is the ratio of their weights, but none of the 
phenomena discussed in this chapter are influenced appreciably 
by their relative sizes, and therefore none of them give any in- 
formation on the subject. Read the footnote to p. 102. 

Exercises. — 1. The weight of 1 1. of a gas at 0° and 760 mm. is 
5.236 g. What is the density referred (a) to air (air = 1) and (6) 
to hydrogen, and (c) what is the molecular weight (pp. 101, 102)? 

2. The density of a gas, referred to air, is 6.7. What is the 
weight of 1 1. (p. 101), and what is the molecular weight (p. 118)? 

3. The molecular weight of a substance is 65. What is the 
density referred to air, and what is the weight of 1 1.? 

4. The chloride of a new element contains 38.11 per cent of 
chlorine and 61.89 per cent of the element. The vapor density of 
the compound referred to air is 12.85. What is the atomic weight 



120 COLLEGE CHEMISTRY 

of the element, so far as investigation of this one substance can 
give it (p. 118)? What is its valence? 

5. If the molecular weight of oxygen were taken as 100, what 
would be the volume of the G.M.V. (p. 101)? What, on the same 
scale, would be the molecular weight of water, and what would be 
the atomic weights of hydrogen and chlorine (pp. 101, 105)? 

6. In future nothing but molecular formulae of free elements 
must be used (p. 111). Write in molecular form ten of the equa- 
tions involving gases which are found in the preceding chapters. 

7. If a new form of oxygen were found, such that one volume 
of it required four volumes of hydrogen to produce water, what 
would be its molecular formula (p. 114)? What would be the 
weight of 22.4 L? 



CHAPTER X 

SOLUTION 

Solutions are so constantly used in chemistry that some 
knowledge of their properties is desirable in order that we may 
employ them intelligently. In what follows, we give a preliminary 
account of some of the simpler facts about solution. 

General Properties of Solutions. — A solid may be dis- 
tributed through a liquid, either by being simply suspended (p. 12) 
in the latter (mixture), or by being dissolved in it (solution). 
Similarly a liquid may be suspended in droplets in another liquid 
(emulsion), as in milk, or it may be dissolved. It is usually easy to 
distinguish between the two cases, for a suspended substance settles 
or separates sooner or later (like the fats in milk — as cream), 
while a dissolved substance shows no such tendency. The cases are 
exceptional where the subdivision of a suspended substance is so 
minute (colloidal suspension, q.v.) , as to make its retention by filter 
paper impossible. If a liquid is opalescent or opaque, then we have 
a case of suspension. A solution is a clear, transparent, perfectly 
homogeneous liquid, in which the dissolved substance seems to 
have been dispersed so completely that the liquid cannot be dis- 
tinguished by the eye from a pure substance. 

There is no limit to the amount of dissipation which may thus 
be produced. A single fragment of potassium permanganate, for 
example, which gives a very deep purple solution in water, may be 
dissolved in a liter or even in twenty liters of water, and the purple 
tinge which it gives to the liquid will still be perfectly perceptible 
in every part of the larger volume. The qualitative characteristics, 
therefore, of solution are absence of settling, homogeneity, and ex- 
tremely minute subdivision of the dissolved substance. 

The Scope of the Word. — The word solution is used for other 
systems than those containing a solid body dissolved in a liquid. 

121 



122 COLLEGE CHEMISTRY 

Thus, liquids also may be dissolved in liquids, as alcohol in water. 
Again, if we warm ordinary water, bubbles of gas appear on the 
sides of the vessel before the water has approached the boiling- 
point. They are found to be gas derived from the air. Agitation 
of any gas with water results in the solution of a large or small 
quantity of the gas, and heat will usually drive the gas out again. 
It appears therefore that solids, liquids, and gases can equally form 
solutions in liquids. 

The absorption of hydrogen by palladium (at all events after a 
certain point), and by iron, takes place in accordance with the 
same laws as the solution of solids in liquids, and the results may be 
described therefore as true solutions. Liquids are in some cases 
absorbed by solids, and homogeneous mixtures of solids with solids 
are perfectly familiar. The sapphire is a solution of a small 
amount of a strongly colored substance, in a large amount of color- 
less aluminium oxide. It may therefore be stated that solution of 
gases, liquids, and solids in solids appears to be possible. 

Limits of Solubility. — The next question which naturally 
occurs to us is as to whether the mingling of two substances in this 
manner has any limits. We find that the results of experiment in 
this direction may be divided into two classes. Some pairs, of 
liquids particularly, may be mixed in any proportions whatever. 
Alcohol and water, and glycerine and water are such pairs. On 
the other hand, at the ordinary laboratory temperature, we can 
scarcely take a fragment of marble (CaCOs) so small that it will 
dissolve completely in 100 c.c. of pure water, for only 0.00013 g. 
dissolves. Under the same conditions any amount of potassium 
chlorate up to about 5 g. will completely disappear after vigorous 
stirring, while 90 g. of ordinary Epsom salts (hydrated magnesium 
sulphate), but not more, may be dissolved in about the same 
amount of water. In fact, most solids may be dissolved in a liquid 
only up to a certain limit, which with different solids may range 
from a scarcely perceptible to a very large amount. No substance 
is absolutely insoluble. But for the sake of brevity we call marble, 
for example, " insoluble" because in most connections it may be so 
considered. 

Chemists have not yet succeeded in explaining these differences 
in solubility, which are often so surprising. Thus, guncotton is 



SOLUTION 123 

soluble in a mixture of alcohol and ether, but not in these liquids 
separately, while cellulose acetate (an allied substance, used in 
making artificial horse-hair) is soluble in these liquids separately, 
but not in the mixture. 

Recognition and Measurement of Solubility, — The only 
method of recognizing with certainty whether a solid is soluble in a 
liquid or not is to filter the mixture and evaporate a few drops of 
the filtrate on a clean watch-glass. For learning how much of the 
body is contained in a given solution, a weighed quantity of the 
solution is evaporated to dryness and the weight of the residue 
determined. 

It must be stated explicitly that in going into solution, as we 
have used the term, a compound dissolves as a whole and, if the 
compound is pure (p. 4), any residue has the same chemical com- 
position as the part which has dissolved. If the residue is a 
different substance, a chemical interaction with the solvent has 
occurred. If, on evaporation, a different substance remains, there 
has also been chemical action. 

Terminology. — In order to describe the relations of the com- 
ponents of a solution, certain conceptions and corresponding 
technical expressions are required. 

It is customary to speak of the substance which, like water in 
most cases, forms the bulk of the solution, as the solvent. To 
express the substance which is dissolved, the word solute is fre- 
quently used, and will be employed when we wish to avoid circum- 
locution. 

The amount of the substance which has been dissolved by a 
given quantity of the solvent is described as the concentration of 
the solution. A solution containing a small proportion of the 
dissolved body is called dilute; it has a small concentration. 
One which contains a larger amount is more concentrated. Very 
"strong" solutions are frequently spoken of simply as concentrated 
solutions. The partial removal of the solvent (as by evaporation) 
is called concentrating, its total removal evaporating to dryness. 

Finally, since there is a limit to the solubility of most substances, 
a solution is described as saturated when the solute has given as 
much material to the solvent as it can. This state is reached after 



124 COLLEGE CHEMISTRY 






prolonged agitation with an excess of the gas, of the liquid, or of 
the finely powdered solid, as the case may be (see pp. 127, 133). 
Other things being equal, the larger the excess, the sooner satura- 
tion is attained. The maximum concentration attainable in this 
way is called the solubility of the substance in a given solvent. 
Note that a saturated solution need not also be a concentrated one. 
It will be very dilute, if the solute is but slightly soluble. 

Units Used in Expressing Concentrations. — The concen- 
trations of solutions, saturated and otherwise, are sometimes 
expressed in physical, and sometimes in chemical, units of weight. 
When physical units are employed, we give the number of grams 
of the solute held in solution by one hundred grams of the solvent. 

The solubilities at 18° of one hundred and forty-two bases and 
salts are given in a table printed inside the cover, at the front of 
this book. 

When chemical units of weight are employed, two different plans 
are possible, and both are in use. Either the equivalent (p. 65) or 
the molecular weights may be taken as a basis of measurement. 
In the former case, the solutions are called normal solutions, and in 
the latter, molar solutions. 

A normal solution contains one gram-equivalent of the solute in one 
liter of solution (not in 1 1. of solvent). The word "equivalent" 
has been used hitherto only of elements, and this application of the 
expression involves an extension of its meaning. An equivalent 
weight of a compound is that amount of it which will interact with 
one equivalent of an element. Thus, a formula-weight of hydro- 
chloric acid HC1 (36.5 g.) is also an equivalent weight, for it con- 
tains 1 g. of hydrogen, and this amount of hydrogen is displaceable 
by one equivalent weight of a metal. A formula-weight of 
sulphuric acid H 2 S0 4 (98 g.), however, contains two equivalents 
of the compound, and a formula-weight of aluminium chloride 
A1C1 3 (133.5 g.) three equivalents. Hence normal solutions of 
these three substances contain respectively 36.5 g. HC1, 49 g. 
H2SO4, and 44.5 g. A1C1 3 per liter of solution. The special 
property of normal solutions is, obviously, that equal volumes of 
two of them contain the exact proportions of the solutes which are 
required for complete interaction. Solutions of this kind are much 
used in quantitative analysis. We frequently use also decinormal 



SOLUTION 125 

or one-tenth normal solutions (0.1 N or iV/10), and seminormal 
(0.5 N or N/2), and six times normal solutions (6 N), and so 
forth. 

A molar solution contains one mole (gram-molecular weight) of the 
solute in one liter of solution (not in 1 1. of solvent). When molec- 
ular formulae (p. 109) are used, this means one gram-formula 
weight per liter. In the cases cited above, the molar solution 
contains 36.5 g. HC1, 98 g. H 2 S0 4 , and 133.5 g. A1C1 3 per liter. 
As will be seen, the concentrations of molar and normal solutions 
are necessarily identical when the radicals are univalent. 

Solution One of the Physical States of Aggregation of 
Matter. — When a solid body dissolves in a liquid, the properties 
of the body undergo a very marked change, which to all appearance 
might be chemical. Yet, besides the ease with which a liquid may 
be removed by evaporation and the solid recovered unchanged, we 
note particularly that the concentration of a saturated solution 
cannot be expressed in terms of integral multiples of the atomic 
weights. We shall see also that the quantity of a solid which a 
liquid may take up varies with the slightest change in temperature. 
Now we do not find the composition of chemical compounds so to 
vary. The solution of a solid may therefore, in general, be likened 
to a change in state of aggregation, similar to the conversion of a 
liquid into a gas or a solid (see p. 126). 

As in other changes of state, so in the process of solution, heat is 
always liberated or absorbed. This is known as heat of solution. 
Thus, one formula-weight of sulphuric acid, in dissolving in a large 
volume of water, liberates 39,170 calories, and one formula-weight 
of ammonium chloride, in dissolving, absorbs 3880 calories. 

As there is danger of confusion arising, we may repeat that a 
compound is homogeneous and its composition is expressible in 
chemical units of weight; a saturated solution is homogeneous but 
its concentration varies with temperature so that atomic weights 
cannot be used to describe its composition; a mixture of two solids, 
or an emulsion of two liquids, is neither homogeneous nor in any 
way definite in composition. 

Molecular View of the State of Solution. — Accepting solu- 
tion as a physical state of aggregation, we may now apply the same 



126 COLLEGE CHEMISTRY 

molecular conceptions to the explanation of the behavior of a sub- 
stance in solution as to matter in the gaseous or liquid states. We 
saw that a solid body, which is ordinarily condensed in a small 
space, can be disseminated by the use of a solvent through a very 
large one. The molecules of the solid become scattered like those 
of a gas or vapor through a much greater volume. We may re- 
gard the dissolved substance as being, practically, in a gaseous or 
quasi-gaseous condition. The molecules are torn apart from one 
another, their cohesion is overcome, and their freedom of motion 
is in a measure restored. It is true that they could not continue 
to occupy this large volume for a moment in the absence of the 
solvent. But we may bring this into relation with the case of a 
vapor by saying that a solid body, like common salt, can evapo- 
rate (i.e., "dissolve") at the ordinary temperature, and occupy a 
large space, only when that space is already filled with a suitable 
liquid. The latter acts as a vehicle for the particles of the solid. 
A volatile liquid, on the contrary, can dissolve in an empty space 
and fill it with its particles without any vehicle being required. 

This conception of the quasi-gaseous condition of a dissolved sub- 
stance would be simply fantastic if it did not lead us to a better 
understanding of the behavior of solutions. It does successfully 
explain many things, such as diffusion, osmotic pressure, and satu- 
ration (see next section). 

It is easy to show that, if we place a quantity of the pure solvent 
(Fig. 55) above a concentrated solution of a substance, and then set 
the arrangement aside, the dissolved body slowly makes its way 
through the liquid (Fig. 56), obliterating the original plane of sepa- 
ration. Eventually the dissolved body scatters itself uniformly 
through the whole. In other words, the particles of the dissolved 
substance exhibit the property of diffusion in the same way as do 
those of gases. 

When the diffusion of a gas is resisted by a suitable partition, we 
find that pressure is exercised upon the walls of the vessel and upon 
the partition. It is possible to show that the particles of a dis- 
solved substance exercise a pressure of a very similar kind. This 
pressure is spoken of as diffusion pressure. This pressure is found 
to be proportional to the concentration of the solution, just as 
gaseous pressure is proportional to the concentration of the gas 
(Boyle's law). 



SOLUTION 



127 



Molecular View of the Process of Solution. — We may 

now apply the same ideas to the process of dissolving, with a view 





Fig. 55. 



Fig. 56. 



more especially to explaining why this process ceases, in spite of the 
presence of excess of the solute, when a certain 
concentration has been reached. If some of 
the material dissolves, why not more? 

Let us suppose that it is the dissolving of 
common salt in water (Fig. 57) which we wish 
to explain in detail. We believe that in the 
solid substance the molecules are closely 
packed together, while in the solution they 
are rather sparsely distributed. The process 
of solution must consist in the loosening of the 
molecules on the surface and their passage into 
the liquid. By diffusion, the free molecules 
will gradually move away from the neighbor- 
hood of the surface of the solid and make room 
for others, and thus, if the system remains 
undisturbed, the liquid will eventually become 
a solution of uniform concentration. If a large 
enough amount of the solid has been provided, the ultimate condi- 
tion will be that of a saturated solution with excess of the solid 




Fig. 57. 



128 COLLEGE CHEMISTRY 

beneath. If we had proper means of measuring it, the tendency 
of the molecules to leave the solid in the presence of a given 
liquid would give the effect of a kind of pressure. This is spoken 
of as solution pressure. 

Now the molecules, after having entered the liquid, move in 
every direction, and consequently some of them will return to the 
solid and attach themselves to it. The frequency with which this 
will occur will be greater as the crowding of particles in the liquid 
increases, so that a stage will eventually be reached at which the 
number of molecules leaving the solid will be no greater than that 
landing upon it in a given time. If the whole of the liquid has 
meanwhile become equally charged with dissolved molecules, there 
will be no chance that the field of liquid immediately round the solid 
will lose them by diffusion, so that a condition of balance or equi- 
librium (p. 89) will have been established: NaCl (solid) ^± NaCl 
(dslvd). The motion of the particles in the liquid produces what 
we have called diffusion pressure; and when the diffusion pressure, 
by the continual increase in the number of dissolved molecules, 
becomes equal to the solution pressure, increase in concentration of 
the solution ceases. It is at this point that we speak of the solution 
as being saturated with respect to the particular substance dissolv- 
ing. The analogy to vapor tension and vapor pressure (p. 88) 
is evident. The foregoing explanation should be compared care- 
fully with that given in the section on h the molecular relations in 
liquids, and in that on equilibrium (pp. 81, 89-90). 

Conditions Affecting the Solubility of a Gas. — When the 
dissolving substance is a gas, led through, or confined above the 
liquid at a definite pressure, the gas dissolves until a state of equi- 
librium between dissolving and emission is reached, for example, 
Oxygen (gas) +± Oxygen (dslvd), and the liquid is then saturated 
with the gas. 

It is found, as the molecular theory would lead us to expect, 
that the concentration of the saturated solution of a gas is propor- 
tional to the pressure at which the gas is supplied (Henry's law). 

This equilibrium, Gas (gaseous) ^ Gas (dslvd), can be reached, 
naturally, from the other direction, namely by starting with a 
solution of the gas and a space above the solution containing, at 
first, none of the gas. The gas leaves the solution until the rates 



SOLUTION 129 

of emission and return become equal. Hence, a gas may be en- 
tirely removed from solution by bubbling a foreign gas through 
the liquid. The bubbles furnish the space to receive the emitted 
gas, and have a large surface, so that the process goes on rapidly. 
The bubbles also escape, and carry with them the emitted gas, so 
that, in this case, there is no re-solution. This is a case of nullify- 
ing one of the two opposed tendencies (p. 90). 

When a mixture of two gases is shaken with a liquid, the 
gases behave independently of each other (Dalton's law, p. 72). 
Each has the same pressure, and therefore the same solubility, 
as it would possess if it alone occupied the whole space above the 
liquid. 

Two Immiscible Solvents: Law of Partition. — An interest- 
ing application of the same ideas may be made to a case which 
occurs very commonly in chemical work. If we shake up a small 
particle of iodine with water, we find that it dissolves slowly, giving 
eventually a saturated but very dilute solution. If now ether in 
sufficient quantity be shaken with the aqueous solution, the greater 
part of the iodine will find its way into the ether, and be contained 
in the brown layer which rises to the top. The process of re- 
moving a substance partially from solution in one solvent and 
securing it in another is called extraction. We find in such cases 
that neither solvent can entirely deprive the other of the whole of 
the dissolved substance, if the latter is soluble in both independ- 
ently. A state of equilibrium is finally reached: I 2 (in Aq) +± I 2 
(in ether). The partition of the substance takes place in propor- 
tion to its solubility in each solvent. It is found that any amount 
of the solute, up to the maximum the system can contain, provided 
this does not involve too high a concentration in either solvent, is 
divided so that the ratio of the concentrations in the two solvents 
is always the same. In the case of iodine divided between water 
and ether, this ratio is about 1 : 200. 

This principle is used in Parke's process (q.v.) for extracting 
silver from molten lead, by means of melted zinc as the second 
solvent. It is employed in separating interesting compounds from 
animal secretions and vegetable extracts, and in purifying such 
compounds. Nicotine from tobacco and cocaine from coca- 
leaves, are secured in this way. 



130 COLLEGE CHEMISTRY 

Influence of Temperature on Solubility. — The quantity of 
a substance' which we can dissolve in a fixed amount of a given 
solvent depends very largely upon the temperature of both. Usu- 
ally the solubility increases with rise in temperature. Measure- 
ments may be made by the method described before (p. 123), using 
excess of the finely powdered solute with different portions of the 
same solvent in vessels kept at different temperatures. The most 
useful way of representing the results is to plot them graphically. 
The diagram (Fig. 58) shows the curves for a few familiar sub- 
stances. The ordinates represent the number of grams of the 
anhydrous compound which is held in solution by 100 g. of water 
in each case. The abscissae represent the temperatures. The con- 
centration for any temperature can be read off at once. Thus, 100 g. 
of water holds 13 g. of potassium nitrate in solution at 0° and 150 g. 
at 73°. The increase in solubility is here enormous. On the 
other hand, the same quantity of water will hold 35.6 g. of sodium 
chloride in solution at 0° and 39 g. at 100°. The difference is shown 
at once when we examine the curves and observe that the line repre- 
senting the solubility of sodium chloride scarcely rises at all between 
0° and 100°, while that of potassium nitrate is extremely steep. 

Cases in which the solubility decreases with rise in temperature 
are less common. The solubility of slaked lime (calcium hydroxide 
Ca(OH) 2 , used to make limewater) is 0.175 g. at 20° and 0.078 g. 
at 100°. Anhydrous sodium sulphate Na 2 S0 4 (Fig. 59, p. 132) is 
another illustration. 

Equilibrium in a Saturated Solution. — Once a solution has 
become saturated, the dissolving substance remains thereafter un- 
changed in amount A greater excess of the solute forces no more 
matter into solution than does a small excess. 

It should be clearly understood, however, that an exchange of 
molecules (p. 128) is still going on between the solid and the solu- 
tion. That this conception is correct may be shown in various 
ways. Thus, if a crystal, the edges or corners of which have been 
broken, is suspended in a saturated solution of the same substance, 
it neither increases nor diminishes in weight. Yet we find that the 
imperfections are removed, and that this takes place by the solu- 
tion of a portion of the substance from the perfect surfaces and its 
deposition upon the imperfect ones. 



SOLUTION 



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132 



COLLEGE CHEMISTRY 



Supersaturated Solutions. — Another very striking proof of 
this may be obtained by saturating water with ordinary Glauber's 
salt (hydrated sodium sulphate Na 2 SO4,10H 2 O) at, say, 30°, at 
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0° 10° 20° 30° 40° 50° 60° 70° 80° 90° 100° 
Temperature 
Fig. 59. 

Na 2 S04 (Fig. 59). The excess of the solid is carefully and com- 
pletely separated from the liquid, and the latter is allowed to cool, 
say to 15°, in a flask loosely stoppered with cotton. 
The solution now contains a much larger amount 
of sodium sulphate (NaaSOJ than at its present 
temperature it could acquire from contact with 
Glauber's salt (13 g. at 15°). Yet in the absence 
of a crystal, with which the above described ex- 
change could take place, no deposition of the dis- 
solved substance begins. The solution may be 
kept indefinitely without alteration. The intro- 
duction, however, of the minutest fragment of 
the decahydrate at once starts the exchange, and 
this is necessarily very much to the disadvantage 
of the solution and the advantage of the crystal: 
10H 2 O + Na 2 S0 4 (dslvd) <=± Na 2 SO 4 ,10H 2 O (solid). The latter 
therefore forms the center of a radiating mass of blade-like proc- 
esses, which sprout with astonishing rapidity through the liquid 
(Fig. 60). 




Fig. 60. 



SOLUTION 133 

Usually the cooling of a concentrated solution leads to the almost 
immediate appearance of crystals spontaneously, and the substance 
is deposited gradually as the temperature falls. But solutions 
of a number of common substances, such as sodium thiosulphate 
(photographer's "hypo") and sodium chlorate, behave like that 
of sodium sulphate. They are said to have a tendency to give 
supersaturated solutions. In general, crystallization can be started 
only by introduction of a specimen of the same substance. The 
smallest particle of the right material floating in the air, if it gains 
admission, will bring about the result. This shows the importance 
of the interchange of molecules, of which we have spoken, for 
establishing equilibrium. 

This phenomenon is similar to the supercooling of water (p. 86), 
which results in crystallization (freezing) when a fragment of ice 
is dropped in. 

Definition of a Saturated Solution: A Warning. — To avoid 
a common misconception, it must be noted that a saturated solution 
must not be defined as one containing all of the solute that it can hold. 
A supersaturated solution holds more. The saturated solution is 
one which contains all of the dissolved solute that it can acquire from 
the undissolved solute. Better still, it is that solution which, when 
placed with excess of the solute, is found to be in equilibrium. 

It must be clearly understood that solution is not a process of 
filling the pores of the liquid. If that were true, approximately 
equal weights of all substances would find accommodation in equal 
volumes of water. The fact is that, for example, 100 c.c. of water 
can dissolve 195 g. of silver fluoride, but only 0.00000035 g. of 
silver iodide, although the space available in the solvent (if there 
is any free space) is the same in both cases. 

The same conclusion is reached when we consider that two forms 
of the same compound have different solubilities. Thus, at 20°, 
Na2SO 4 ,10H 2 O can give about 18 g. of NasSC^ to 100 c.c. of water 
(Fig. 59). But anhydrous sodium sulphate Na2S0 4 at 20° gives 
59 g. to the same amount of water (read up to dotted line, Fig. 59). 
Note that in the diagram (Fig. 59) the solubility curve of Na^SC^, 
10H 2 O comes to an end at 32.4°. At this temperature the solid 
melts and decomposes, so that measurements with this solid be- 
yond that temperature are impossible. 



134 COLLEGE CHEMISTRY 

Influences of the Solute Upon the Solvent. — These in- 
fluences are of two classes. In one of these classes, equal num- 
bers of dissolved molecules of different substances produce the 
same amount of change. The effect appears, therefore, to be 
largely due to mechanical causes. Of this nature are the lowering 
of the freezing-point of the liquid, the lowering of its vapor tension 
and the raising of its boiling-point, and the value of the osmotic 
pressure (see below). 

In the other class, the effect varies with the substance dissolved. 
The changes in volume (see below) belong to this class. 

Freezing-Points of Solutions: Freezing Mixtures, — The 

freezing of a dilute solution consists, usually, in the crystalliza- 
tion of some of the pure solvent only. The presence of a dissolved 
body tends to prevent this freezing, and so solutions can be frozen 
only at temperatures below those at which the pure solvent would 
freeze. Thus, one gram-molecular weight of any* substance, 
such as sugar (342 g.) or alcohol (46 g.), dissolved in 1000 c.c. 
(1 liter) of water, will cause the water to freeze at —1.86° instead 
of at 0°. 

This explains why sea water is much less often frozen in cold 
weather than is fresh water. It should be noted, also, that the ice 
formed in salt water is free from salt. 

This fact likewise explains why salt thrown on ice causes the 
latter to melt. A saturated solution of salt does not freeze until 
cooled to —21° ( — 6° F.), and it then gives a mixture of pure ice 
and pure salt crystals. Hence, ice and salt cannot permanently 
exist together above —21°. Below —6° F., salt will no longer melt 
ice. A mixture of ice and salt, giving a temperature of —6° F., 
is called a freezing mixture, and is used in freezing ice cream and 
ices. 

Molecular weights of non-volatile substances can be measured by 
simply finding out what weight of the substance, in 1000 c.c. of 
water, is required to lower the freezing point from 0° to —1.86°. 

The Vapor Tension of Solutions: Deliquescence. — A 

solute, which is itself noifc-volatile, tends to diminish the vapor 
tension of the solvent. It hinders the emission of vapor. 
* For important exceptions, see Chap. XVI. 



SOLUTION 135 

If the substance is very soluble, and the solution highly concen- 
trated, the lowering in the vapor tension will be considerable. In 
fact, the solution may give a vapor pressure of water less than that 
commonly present in the atmosphere. Such a solution, placed in 
an open vessel, will not evaporate. On the contrary, vapor from 
the air will enter it, and it will increase in bulk. For this reason, 
crystals of very soluble substances are usually moist and, when 
exposed to the air, acquire water from the latter and dissolve in 
this water. This behavior is called deliquescence, and is exhibited, 
for example, by the hydrate of calcium chloride CaCl 2 ,2H 2 0, 
which is consequently used for drying gases. Magnesium chloride 
MgCl 2 , present as an impurity in common salt, causes the latter to 
become moist in damp weather. 

The principle involved will become clear if we imagine two 
vessels, one containing pure water and one an aqueous solution, 
to be placed on a glass plate and covered by a bell jar (Fig. 61). 
Each liquid exchanges water molecules with the moist air in the 
jar, but the solution gives off water more 
feebly than does the pure water. The 
result is that the latter can produce a pres- 
sure of water vapor higher than that which 
would be in equilibrium with the solution. 
The solution, therefore, receives continu- 
ously more molecules than it emits, and 
increases in volume. The pure water thus 
gradually passes through the vapor state 

into the solution until it is all gone. If sufficient water was 
present, the process would go on until the solution became infinitely 
dilute. 

Boiling-Points of Solutions. — Since the solute interferes 
with the emission of the vapor of the solvent, it naturally makes 
the solution more difficult to boil. It raises the boiling-point. 
Thus one gram-molecular weight of sugar (342 g.) or of glycerine 
(92 g.), dissolved in 1000 c.c. of water, will elevate the boiling 
point from 100° to 100.52° (for exceptions, see Chap. XVI). 

The Laws of Osmotic Pressure. — We have seen that a dis- 
solved substance exercises a pressure, called osmotic pressure, which 




V I T 



136 



COLLEGE CHEMISTRY 



is proportional to the concentration of the solute in the solvent. 
This pressure shows itself when a membrane opposes the diffusion 
of the solute from the solution into a layer of pure solvent. The 
phenomena can be illustrated by using a diffusion shell, of test- 
tube form, attached at the lower end of a glass tube (Fig. 62). 
We may place the solution {e.g., of sugar) inside the shell, and 
pure water outside. The material of the 
shell is such that the water can pass through 
it in either direction, but the molecules of 
the dissolved substance cannot do so. Such 
a membrane is called semi-permeable. 

The facts observed with this arrangement 
are as follows: (1) The pure solvent passes 
into the solution. Thus, if the solution is 
inside the shell, water enters through the wall 
of the shell, and the liquid therefore rises in 
the tube. If the solution is outside, water 
passes out of the shell, and the liquid in the 
tube falls. (2) If solutions of different con- 
centrations are used inside and outside, the 
solvent passes from the more dilute into the 
more concentrated solution, so that the tend- 
ency is to dilute the latter and concentrate 
the former until both have the same con- 
centration. (3) The entrance of the solvent 
can be prevented by the application of pressure 
to the surface of the liquid in the tube. With 
1 per cent sugar solution inside and water 
outside (at 15°), a pressure equal to 500 mm. 
of mercury (0.66 atmos.) per square centimeter of the surface 
of the membrane is required to stop the entrance of the sol- 
vent. This, therefore, is the value of the so-called osmotic pres- 
sure. Since the entrance of the solvent is due to the dissolved 
substance, and the solvent is really drawn forcefully into the 
solution, it might be more appropriate to call the force osmotic suc- 
tion. Whatever it is named, however, it is real, and its value can 
be measured. (4) The value of the pressure (or suction) increases 
in proportion to the absolute temperature, just like the pressure 
of a gas. (5) The value of the pressure (or suction) is also propor- 




Fig. 62. 



SOLUTION 137 

tional to the concentration of the solute in the solution. Thus, 2 per 
cent sugar gave (Pfeffer) 1016 mm. pressure, 4 per cent sugar 2082 
mm. (6) When different solutes are compared, it is found that, 
at the same temperature, equal numbers of molecules of the solutes 
dissolved in equal volumes of the solvent give almost equal osmotic 
pressures (or suctions). Thus, one mole (342 g.) of sugar C12H22O11 
and one mole (74 g.) of methyl acetate CH3CO2CH3, which, in 
spite of the great difference in weight, contain equal numbers of 
molecules, when dissolved separately, each in ten liters of water, 
give in both cases 2.42 atmospheres osmotic pressure (or suction). 
(7) Finally, the osmotic pressure (or suction) exercised by a certain 
quantity of a substance in dilute solution is identical in value with 
the gaseous pressure which the same quantity of the same substance 
would exercise if it were contained as a gas in the same volume at the 
same temperature. Thus 44 g. of carbon dioxide, as a gas, filling 
22.4 liters (the G.M.V.) at 0° exercises one atmosphere gaseous 
pressure. When the cube is filled with water, and the gas is 
thus dissolved, the osmotic pressure of the solution is one 
atmosphere. 

These facts apply to substances which are not acids, bases, 
or salts. W^e shall find later (see Chap. XVI) that the osmotic 
pressures of the members of these three classes of substances are 
frequently abnormally high, but that the abnormality is easily 
explained. 

Osmotic pressure (or suction) is a subject of great interest in 
connection with the physiology of plants and animals. It aids in 
explaining why a withered flower, containing a solution in its cells, 
revives when placed in pure water. The latter enters through the 
walls of the cells, and the pressure thus produced distends the 
structure and stiffens it. Similarly, the ascent of the water from 
the soil into the roots and through the stem of a growing plant is 
explained. In the animal body also, osmosis plays a large part. 

Measurements of osmotic pressures cannot be made accurately 
with a diffusion shell, because the solute is able to some extent to 
pass through the material. Then, too, such a shell can be used 
only with dilute solutions, because it will not withstand high 
pressures. In making accurate measurements, a cell of porous 
porcelain is used, and the pores are filled with a gelatinous precipi- 
tate of cupric ferrocyanide Cu 2 Fe(CN) 6 {q.v.). 



138 COLLEGE CHEMISTRY 

Densities of Solutions. — The density or specific gravity of 
a solution is usually greater than that of water and, in each case, 
varies with the concentration. For commercial purposes, the 
concentrations of solutions are commonly defined by the specific 
gravity. Thus, we purchase ammonium hydroxide solution of 
"0.88 sp. gr.," meaning 35 per cent of ammonia, or sulphuric acid 
of "1.84 sp. gr.," meaning 94.8 per cent of the acid. 

The commonly greater density of a solution is utilized in making 
solutions in chemical factories. Shaking several tons of the 
mixture is out of the question, and stirring costs money. If the 
solid is placed in the bottom of the tank, under water, a. saturated 
solution is formed in the lowest layer of the water, and passage of 
the dissolving substance into the upper layers, by diffusion, would 
take months or years. Hence most of the solid would remain un- 
dissolved (Fig. 57, p. 127). But when the solid is placed on a 
shelf near the surface of the water, the solution sinks through the 
water, fresh water rises to the shelf, and a circulation is started. 
This results in the dissolving of the whole material in a surprisingly 
short time, with no expenditure of labor whatever. 

Changes in Volume upon Solution. — The erratic and, at 
present, unexplained changes in volume which occur when a sub- 
stance is dissolved, seem to indicate that the process is less simple 
than we have thus far admitted, and that chemical changes occur 
during the process. Thus, when 250 g. of common salt are dis- 
solved in 1 liter of water (= 1000 c.c. = 1000 g.), which gives a 
20 per cent solution, the volume of the solution is only 1086 c.c. 
Since the 250 g. of salt occupied 116 c.c. before being dissolved, a 
shrinkage of 1116 — 1086 or 30 c.c. accompanied the process of solu- 
tion. On the other hand, 214 g. of ammonium chloride (volume 
142.5 c.c.) and 843.5 c.c. of water, have a total volume of 986 c.c, 
but when dissolved give 1000 c.c. of solution. Here there is an 
expansion of 14 c.c. Table sugar, however, dissolves in water 
with almost no change in volume. 

Is Solution a Physical or a Chemical Change? — These 
phenomena are, in part, accounted for by the fact that water is not 
a single substance, but a mixture. It is largely composed of 
dihydrol (H 2 0) 2 , with much trihydrol (H 2 0) 3 near to 0° and in- 
creasing quantities of monohydrol H 2 at higher temperatures. 



SOLUTION 139 

When any substance is dissolved in considerable amount in water, 
the equilibrium amongst these three kinds of molecules is dis- 
turbed, and then proportions change: 

2(H 2 0) 3 *=> 3(H 2 0) 2 ^ 6H 2 0. 

Now, equal weights of these three kinds of water occupy different 
volumes, and hence solution is accompanied by changes in the 
volume of the water. The same condition in water explains the 
point of maximum density (4°). The change from (H 2 0)3 to 
(H 2 0) 2 , which proceeds as the temperature rises from 0° to 4°, is 
accompanied by a shrinkage, because dihydrol has the higher 
density. Beyond 4°, the usual expansion with rising temperature 
prevails. 

There is also evidence that many dissolved bodies form unstable 
compounds with water, although we have not as yet definite in- 
formation about these compounds. 

Dissolving in water is, therefore, partly a chemical and only 
partly a physical process — a part of the water is always affected, 
and a part or all of the solute may go into combination. 

Exercises, — 1. Give other examples of limited solubility in 
various solvents (p. 122). 

2. What weights of phosphoric acid (p. 94) and of sodium 
hydroxide, respectively, are required to make 1 liter of a normal 
solution? 

3. Express the concentrations of solutions of ammonium 
chloride, saturated at 0° (sp. gr. 1.076), and of potassium sulphate 
K 2 S0 4 , saturated at 10° (sp. gr. 1.083), in terms of a normal solu- 
tion (p. 124). 

4. Express the concentration of a five per cent aqueous solution 
of phosphoric acid (sp. gr. 1.027), in terms of a normal and a molar 
solution, respectively. 

5. Explain why, (a) pulverization and, (6) agitation hasten the 
dissolving of a solid (c/. pp. 331, 398). 

6. Read from the curves (p. 131) the solubilities of potassium 
nitrate at 15°, of potassium chloride at 30°, of potassium chlorate 
at 45°. What are the relative rates at which the solubilities of 
these salts increase with rise in temperature? 

7. If 5 g. of a substance, dissolved in 1000 c.c. of water, give a 



140 COLLEGE CHEMISTRY 

solution freezing at —0.2°, what is the molecular weight of the 
substance? 

8. At what point in a tank of water should you introduce 
ammonia gas, in order, with the least effort, to saturate the water? 
The sp. gr. of the saturated solution is 0.88. 

9. 6 grams of a substance when dissolved in 200 c.c. of water 
give a boiling-point of 102.6°. What is the molecular weight of 
the substance? 

10. 1.6 grams of naphthalene GoH 8 when dissolved in 25 g. of 
benzene (freezing-point 5.48°) gives a solution which freezes at 
3.03°. When 2.44 grams of another substance are dissolved in 
the same amount of benzene, the solution freezes at 3.52°. What 
is the molecular weight of the latter substance? 

11. The elevation of the boiling-point in the above solution of 
naphthalene is 1.285°. What elevation of the boiling-point is pro- 
duced in the second solution? 



CHAPTER XI 



HYDROGEN CHLORIDE. CALCULATIONS 



We have had occasion several times to mention common salt, 
or sodium chloride NaCl. This is one of the most familiar chemi- 
cal substances. Large quantities of it are used in the household, 
in cooking and in making freezing mixtures. Still larger amounts 
are consumed in manufacturing washing soda, caustic soda, and 
soap, for all of which it furnishes the necessary sodium. It is 
used also in preserving fish and other foods. It supplies the 
chlorine used in bleaching and in the steriliza- 
tion of city waters. We shall consider it first 
as a means of making other compounds of 
chlorine. 



Preparation of Hydrogen Chloride HCl 
from Salt, — When some concentrated sul- 
phuric acid is poured upon sodium chloride, a 
vigorous effervescence is noticed. This shows 
that bubbles of a gas are forming upon the 
salt crystals and are rising through the acid 
and bursting. If the salt be placed in a 
flask (Fig. 63), the acid can be allowed to enter 
from time to time through the funnel. When 
the air has been displaced from the flask, the gas issues from 
the delivery tube. If the correct proportion of the acid is used, 
and only a gentle heat is applied, all that remains in the flask is 
a white solid, sodium-hydrogen sulphate (sodium bisulphate) 
NaHS0 4 : 

NaCl + H 2 S0 4 -> NaHS0 4 + HCl t .* (1) 

* The arrow directed downwards indicates elimination of a substance by 
precipitation; that directed upwards, escape as a gas or solution of a solid. 

141 




Fig. 63. 



142 COLLEGE CHEMISTRY 

The gas is extremely soluble in water and, being heavier than 
air, may be collected by upward displacement of the air in a jar. 

The action described is the one which occurs in the laboratory. 
When a double proportion of salt and a high temperature are used, 
a second action occurs: 

NaCl + NaHS0 4 -> Na 2 S0 4 + HC1 T 

and sodium sulphate Na 2 S0 4 remains. In Europe this action is 
employed, with furnace heat, in manufacturing sodium sulphate, 
from which sodium carbonate is afterwards prepared. The hydro- 
gen chloride passes into a tower, down which water trickles over 
lumps of coke, and is dissolved. The aqueous solution is called 
hydrochloric acid or, in commerce, muriatic acid (Lat., brine 
acid) . 

Hydrogen Chloride from Other Chlorides and Other 
Acids. — The chlorides of other metals could be substituted for 
sodium chloride in this action, and all the more soluble ones would 
give hydrogen chloride freely. Other chlorides are all more ex- 
pensive, however, than is common salt. 

All acids contain the necessary hydrogen radical, and might 
offer it in exchange for the sodium in the salt, yet in practice no 
other acid works so well as does sulphuric acid. Concentrated 
phosphoric acid H 3 P04,Aq acts more slowly, giving primary 
sodium phosphate: 

NaCl + H 3 P0 4 -> NaH 2 P0 4 + HC1 1 . 

The Molecular View of the Interaction of Sulphuric Acid 
and Salt, — One who has used the above-described methods for 
making hydrogen chloride without reflection would not realize the 
complexity of the machinery by which the result is achieved. The 
means are apparently very simple. Yet the mechanical features 
of this experiment, when laid bare, are extremely curious and in- 
teresting. A single fact will show the possibilities which are 
concealed in it. 

If we take a saturated solution of sodium-hydrogen sulphate in 
water and add to it a concentrated solution of hydrogen chloride in 
water (concentrated hydrochloric acid), we shall perceive at once 



HYDROGEN CHLORIDE 143 

the formation of a copious precipitate. This is composed entirely 
of minute cubes of sodium chloride: 

NaHS0 4 + HC1 -> H 2 S0 4 + NaCl I* (2) 

Now this action is nothing less than the precise reverse of (1), yet 
it proceeds with equal success. In fact, this chemical interaction 
is not only reversible (pp. 93, 95)', but can be carried to comple- 
tion in either direction. It is only in presence of a large amount 
of water that it stops midway in its career and is valueless for 
securing a complete transformation in either direction: 

NaHS0 4 + HC1 <± H 2 S0 4 + NaCl. 

In an action which is reversible, if the products remain as per- 
fectly mixed and accessible to each other as were the initial sub- 
stances, their interaction will continually undo a part of the work 
of the forward direction of the change. Hence, in such a case the 
reaction must, and does, come to a standstill while as yet only 
partly accomplished; but this was not the case with actions 
(1) and (2). Let us examine the means by which the premature 
cessation of each was avoided. 

In equation (1) the salt dissolved to some extent in the sulphuric 
acid, NaCl (solid) ±± NaCl (dslvd), and so, by contact of the two 
kinds of molecules, the products were formed. On the other hand, 
the hydrogen chloride, being insoluble in sulphuric acid, escaped as 
fast as it was formed: HC1 (dslvd) ^ HC1 (gas). Hence, in that 
case, almost no reverse action was possible, and the double decom- 
position went on to completion. With all the sodium-hydrogen 
sulphate in the bottom of the flask, and most of the hydrogen 
chloride in the space above, the two products might as well have 
been in separate vessels so far as any efficient re-interaction was 
concerned. This plan, in which water is purposely excluded, forms 
therefore the method of making hydrogen chloride. 

In equation (2), on the other hand, the hydrogen chloride was 
taken in aqueous solution, and was mixed with a strong solution of 
sodium bisulphate. The acid was, therefore, kept permanently in 
full contact with the sodium bisulphate. It had in this case, every 
opportunity to interact with the latter and no chance of escape. 
Every molecule of each ingredient could reach every molecule of 
* See footnote to p. 141. 



144 COLLEGE CHEMISTRY 

the other with equal ease. Furthermore, the sodium chloride, 
produced as a result of their activity, is not very soluble in con- 
centrated hydrochloric acid (far less so than in water), and so it 
came out as a precipitate: NaCl (dslvd)^NaCl (solid). But 
this was almost the same as if it had gone off as a gas. It meant 
that the greater part of the salt was in the solid form. It was in 
a state of fine powder, it is true. But, in the molecular point of 
view, the smallest particle of a powder contains millions of mole- 
cules, and most of these are necessarily buried in the interior of a 
particle. Thus, the sodium chloride was no longer able to interact 
effectively molecule to molecule with the other product, the sul- 
phuric acid. Hence, there was little reverse action to impede the 
progress of the primary one. Thus (2) is nearly as perfect a way 
of liberating sulphuric acid as (1) is of liberating hydrogen 
chloride. 

This discussion is given to illustrate the displacement of a chemi- 
cal equilibrium, and to explain the method of preparing hydrogen 
chloride. It also throws an interesting light on chemical affinity, 
however. Considering action (1), by itself, we might reason that 
the hydrogen chloride was formed because the affinity of the hydro- 
gen (H) for chlorine (CI) was greater than for the sulphate radical 
(S0 4 ). But, if we did so, then in action (2) we should be compelled 
to reason similarly that the preponderance of affinity was just the 
opposite. In point of fact, no conclusion about relative affinity 
can be drawn from these actions. The effects of affinity are here 
entirely subordinated by the effects of a purely mechanical ar- 
rangement, depending on solubility. When the activities of the 
acids are properly compared, hydrochloric acid is found to be 
considerably more active than sulphuric acid. 

Physical Properties. — Hydrogen chloride is a colorless gas, 
which produces a suffocating effect when inhaled. 

Density (H = 1), 18.23. Crit. temp., +52°. 

Weight of 22.4 L, 36.73 g. Boiling-point (liq.), -83.7°. 

Sol'ty in Aq (0°), 50,300 vols, in 100. Melting-point (solid), -110°. 

The gas is one-fourth heavier than air. On account of its great 
solubility, when it streams into the air it condenses atmospheric 
moisture into a fog (of drops of hydrochloric acid). The extreme 



HYDROGEN CHLORIDE 



145 



solubility may be shown by filling a dry flask (Fig. 64) with the 
gas. The " dropper" contains water, and is closed at the tip 
with soft wax. A drop of water, expelled by pinching the "drop- 
per/' dissolves so much of the gas that the water is forced in by 
atmospheric pressure, like a fountain, through the 
longer tube. 

Both in the gaseous and liquefied states it is a non- 
conductor of electricity. Its heat of solution is 
17,400 calories (p. 85). On account of its high 
concentration, the saturated, aqueous solution may 
be looked upon as a mixture of liquefied hydrogen 
chloride and water. 

When the concentrated aqueous solution is heated, it 
is the gas and not the water which is driven out, for 
the most part. When the concentration has been 
reduced to 20.2 per cent, the rest of the mixture 
distils unchanged at 110°. This occurs because, at 
this concentration, the gas is carried off in the bubbles 
of steam in the same proportion in which it is present in the 
liquid. If a dilute solution is used, water is the chief product of 
distillation (about 100°), but gradually the boiling-point rises and, 
when the concentration has reached 20.2 per cent once more, 
the same hydrochloric acid of constant boiling-point (110° at 760 
mm.), as it is called, forms the residue. 




Fig. 64. 



Chemical Properties. — Hydrogen chloride is extremely stable, 
as we might expect from the vigor with which the elements of which 
it is composed combine (see p. 160) . On being heated to a tempera- 
ture of 1800°, however, it begins to dissociate into its constituents. 

In the chemical point of view, it is on the whole rather an indif- 
ferent substance. Hydrogen chloride (the gas) has no action upon 
any of the non-metals, such as phosphorus, carbon, sulphur, etc. 
Many of the metals, however, particularly the more active ones, 
such as potassium, sodium, and magnesium, decompose it. Hy- 
drogen is set free, and the chloride of the metal is formed. The 
equation representing the weights is K + HO — -> KC1 + H. But 
the molecular formula (p. Ill) of hydrogen is H 2 , hence the cor- 
rect equation is: 

2K + 2HC1 -» 2KC1 + H 2 . 



146 COLLEGE CHEMISTRY 






Hydrogen chloride unites directly with ammonia gas to form a 
cloud of solid particles of ammonium chloride (HC1 + NH 3 — > 
NH^Cl). 

Chemical Properties of Hydrochloric Acid. — The solution 
of hydrogen chloride in water is an entirely different substance in 
its chemical behavior from hydrogen chloride. It is strongly acid, 
turning blue litmus red. The gas and liquefied gas have no such 
property. The solution conducts electricity very well, and is de- 
composed in the process (p. 55), giving hydrogen at the negative 
wire and chlorine at the positive wire: 

2HC1-+H 2 (neg. wire) + Cl 2 (pos. wire). 

The gas and the liquefied gas are practically nonconductors. 

The metals preceding hydrogen in the order of activity (p. 60), 
when introduced into hydrochloric acid, displace the hydrogen 
(p. 55), and form the chloride of the metal. In the case of zinc 
the action was represented by the equation: 

Zn + 2HCl->ZnCl 2 + H 2 . 

The aqueous solution of hydrogen chloride interacts rapidly with 
most oxides and hydroxides of metals, as, for example, those of zinc: 

ZnO + 2HC1 -> ZnCl 2 + H 2 0, 
Zn(OH) 2 + 2HC1 -» ZnCl 2 + 2H 2 0. 

Here no free hydrogen is obtained, since the oxygen in the oxide, 
and the hydroxyl in the hydroxide, unite with it to form water. 
In each case, however, the chloride of the metal is obtained. It 
may be noted, in passing, that all acids behave in a similar manner 
towards oxides and hydroxides of metals, giving water and a com- 
pound corresponding to the chloride. Dilute sulphuric acid, for 
example, gives sulphates. 

Modes of Preparing Chlorides, — In the preceding section 
three kinds of actions, each constituting a different mode of pre- 
paring chlorides, have been mentioned incidentally. There are 
two others. The simplest is the direct union of the element with 
chlorine (Zn + Cl 2 — » ZnCl 2 ) . The other method is illustrated 
in the case of the precipitation of silver chloride by adding asolu- 



HYDROGEN CHLORIDE 147 

tion of a chloride to a solution of silver nitrate. Here the forma- 
tion of the chloride occurs by exchange of another radical (p. 53) 
for the chloride radical: 

AgN0 3 H- NaCl -» AgCl J + NaN0 3 . 

The insoluble chlorides (see p. 164) can be made conveniently by 
this plan. The formation of the precipitates, for example that of 
silver chloride, is used as a test for the presence of a soluble chlo- 
ride in the solution. 

Uses of Hydrochloric Acid, — This substance is used, in 
Europe, as a commercial source of chlorine. It is employed in 
cleaning metals, and in the manufacture of chlorides of metals. 
It is an important component of the gastric juice of the stomach, 
although the proportion is only about 1 part in 500. 

Precipitation. — When two soluble substances are dissolved, 
separately, and the solutions are mixed, chemical interaction fre- 
quently occurs, as in the case of salt and silver nitrate (see also 
p. 143). If one of the products is insoluble, then a supersaturated 
solution of this product is at once produced. As a rule, this sub- 
stance almost immediately becomes visible as a fine powder, called 
a precipitate, suspended in the liquid. 

The insoluble product can often be recognized by its physical 
appearance, and so this sort of action is frequently used as a 
test for one of the original substances. Thus many precipitates 
have a distinctive color. Again, precipitates which are colorless, 
or have the same color, differ in appearance, and are described as 
gelatinous, curdy, pulverulent, or crystalline. In the first two cases, 
the precipitation is so sudden that there is no time for crystals to 
be formed, and the product is amorphous (Gk., without form). 
Thus silver chloride is curdy, and precipitated sodium chloride 
(p. 143) is crystalline. 

Fourth Variety of Chemical Change: Double Decompo- 
sition. — In this chapter we encounter for the first time the 
fourth variety of chemical change. Upon examining the equa- 
tion for the action of sodium chloride and silver nitrate, we see 
that the silver nitrate decomposed into its radicals (Ag) and (N0 3 ). 



148 COLLEGE CHEMISTRY 

The sodium chloride also decomposed into its radicals (Na) and 
(CI). The (Ag) then united with the (CI) and the (Na) with the 
(N0 3 ). 

AgN0 3 + NaCl -> AgCl -f NaN0 3 . 

Since both of the original substances decomposed, this is called a 
double decomposition. An exchange of radicals occurred. 

The action by which hydrogen chloride was prepared (p. 142) 
belonged to the same class: 

NaCl + HHSO4 -> NaHS0 4 + HC1. 

Double decompositions involving acids, bases, and salts are all 
reversible reactions. The fact that many of them proceed, never- 
theless, to practical completion has already been explained at length 
(pp. 142-144). 

The Varieties of Chemical Change, — Most chemical 
changes belong to one of the four varieties: 

1. Combination, e.g., Fe -j- S — » FeS. 

2. Decomposition, e.g., 2KC10 3 -> 2KC1 + 30 2 . 

3. Displacement, e.g., Zn + 2HC1 -► H 2 + ZnCl 2 . 

4. Double Decomposition, e.g., AgN0 3 + HC1 -> AgCl + HN0 3 . 
In the first, 2 (or more) substances give 1 substance. 

In the second, 1 substance gives 2 (or more) substances. 

In the third, 1 element and 1 compound give 1 element and 1 
compound. 

In the fourth, 2 compounds give 2 compounds. 

Occasionally, one compound gives one (different) compound, a 
change called internal rearrangement. Nearly all chemical changes, 
so far as their mechanism is concerned, can be classified under 
one or other of these five kinds. 

A dissociation (p. 93) is both a combination and a decomposi- 
tion, because it is reversible. For example: 

2H 2 0*±2H 2 + 2 . 

Electrolysis is decomposition by an electric current. 

The foregoing varieties of chemical action are general, and not 
limited to any classes of elements. Oxidation (p. 36) and reduc- 
tion (p. 37) are so limited. Thus combination with oxygen is 
oxidation, while combination with hydrogen is reduction. A more 



CALCULATIONS 149 

complete discussion of action of these classes will be given later 
(see Chapter XXIII). 

The reader should classify each action mentioned in the text, 
and so become familiar with the chemical point of view which this 
classification represents. 

Salts. — We have seen that an acid contains hydrogen H as 
a radical (p. 52), and a base contains the radical hydroxyl OH 
(p. 94) . The name salts is given to the class of substances which 
contain a positive and a negative radical, neither of which is hydro- 
gen nor hydroxyl. For example, NaCl, Na2S0 4 , AgNOs are the 
formulae of salts. Salts are so named because they resemble 
common salt in having two radicals, and entering readily into 
double decomposition. 

Sodium-hydrogen sulphate NaHS04 is classed as an acid salt, 
because it has a positive and a negative radical, and a hydrogen 
radical in addition. 

Calculations 

Familiarity with the interpretation of molecular equations is 
best obtained by making simple calculations based upon their 
common uses in chemistry. 

Weights. — When a problem in regard to weights of material 
used or produced in a given action is to be solved, the molecular 
equation is to be written and the weights inserted beneath the 
formulae. The mode of calculation has been described already 
(pp.67, 116). 

Weights and Volumes. — When a problem involving weights 
and volumes is to be solved, the molecular equation is to be written, 
and both the weights and volumes are to be inserted. Note, how- 
ever, that only the volumes of the substances in the gaseous condi- 
tion are considered. 

For example, what volume of oxygen is obtained from 60 g. of 
potassium chlorate? The molecular equation, made as already 
described (p. 116), together with the full interpretation, are as 
follows : 




150 COLLEGE CHEMISTRY 

2KC10 3 -» 2KC1 

Weights- p (39.1 + 35.46 + 48) 2(39.1 + 35.46 ) 
WEIGHTS - [ 24517 14917^ 

Volumes: 

Observe that no volumes are given under the chlorate and chlo- 
ride of potassium. This is because their volumes in the gaseous 
condition can be of no practical use, since they are solids which are 
melted, but not vaporized during this, or any action in which we 
employ them. Now, as to the problem in hand, it is concerned 
with a weight of potassium chlorate and a volume of oxygen. 
Reading from the equation, our information on these points is 
that 245.1 g. of potassium chlorate give 67.2 liters (observe that 
the coefficients are used, as well as the molecular weights, in these 
numbers) of oxygen at 0° and 760 mm., and the question is: What 
volume will 60 g. give? By proportion, 245.1 g. : 60 g. : : 67.2 1. : x 1., 
where x = 16.45 liters. If a different temperature and pressure 
had been specified, either the volume in the equation, or the an- 
swer, would have had to be converted, by rule, to the given condi- 
tions. 

It saves time not to write out, as above, the whole interpreta- 
tion, but only the parts required. For example, if the question 
is: What volume of chlorine is needed to give 25 g. of aluminium 
chloride? we may, if we choose, omit all the data excepting the 
volume of the chlorine and the weight of the aluminium chloride, 
thus: 

2A1 + 3C1 2 -> 2AICI3 
3 X 22.4 1. 2 X 133.5 g. 

The volume of chlorine required is 25 X 3 X 22.4 -J- (2 X 133.5) 
liters. These illustrations show the method of calculating actual 
volumes (see Exercises 1, 2). 

Relative Volumes Alone, — If the question concerns relative 
volumes only, then it is simplest to use the interpretation- of the 
equation in terms of molecules. For example: What relative 
volumes of hydrogen chloride and oxygen are required in Deacon's 
process (see p. 155)? The molecular equation is 

4HCl + 2 ->2H 2 + 2Cl 2 . 
Molecules: 4 12 2 



CALCULATIONS 151 

Since equal numbers of molecules of gases occupy equal volumes, 
the proportion 4 molecules of hydrogen chloride to 1 molecule of 
oxygen shows the ratio to be 4 : 1 by volume. Similarly, every 4 
molecules of hydrogen chloride give 2 molecules of chlorine, so that 
the ratio of these substances by volume is 4: 2, or 2 : 1. 

In regard to the water, since that is not a gas at common tem- 
peratures, the question, if asked, must be more specific : What are 
the relative volumes of steam and chlorine in the product, as com- 
monly delivered by this action at 400°? It is 2:2, or 1:1. 
What are the relative volumes of water and chlorine, after the 
products have cooled to room temperature? The water is no 
longer a gas, so that it occupies, relatively, almost no volume.* 

What is the total volume-change in the foregoing action above 
100°? It is a change from 5 molecules to 4. The volume changes 
in the same ratio. But at 0° the volume-change is from 5 volumes 
to 2, for the water does not appreciably add to the volume of the 
products (see Exercises 3, 4). 

Relative Volumes, Again. — When we know the molecular 
formulae of the single substances concerned in an action, the equa- 
tion can be made, and the relative volumes determined, without 
actual measurement. For example : What volume-change will be 
observed when a mixture of carbon monoxide and oxygen has ex- 
ploded, and the temperature has once more reached that of the 
room? The molecular formulae are CO, 2 , and C0 2 . The equa- 
tion representing the weights is CO + O — > C0 2 . The molecule 
of oxygen, however, being 2 , we cannot employ less than this 
quantity in a molecular equation, so that the equation becomes: 

2CO + 2 ->2C0 2 . 

Three molecules, therefore, give two, throughout the whole mass, 
and therefore three volumes will become two, if the pressure and 
temperature are the same at the beginning and end of the action. 

* Of course if an exact answer must be given, it can be given. But for 
this we require the weight and specific gravity of the product. Thus, 2H 2 
represents 2 X 18 g. of water. The sp. gr. of water is 1. Therefore the 
volume of water formed is 36 c.c. The volume of 2C1 2 is 2 X 22.4, or 44.8 
liters at 0°. The ratio of water to chlorine by volume at 0° is therefore 36: 
44,800. But, aa a rule, we simply give the volumes of solids and liquids as 
zero, compared with those of the gases concerned in the same action. 



152 COLLEGE CHEMISTRY 

If we remember that all volatile compounds of carbon and 
hydrogen burn to form water and carbon dioxide, the molecular 
equation for any such combustion may easily be made, and the 
volumes of all the materials ascertained. When water is a product, 
only its volume as steam is given by the equation (see Exercises 
4, 5). 

Relative Densities of Gases. — Knowing by heart the molec- 
ular formulae of gaseous substances, as we must know them for 
many purposes, it is unnecessary to burden our minds with other 
data in regard to the relative weights of gases. Is hydrogen chloride 
(HC1) heavier or lighter than carbon dioxide (C0 2 )? These for- 
mulae represent the weights of equal volumes (22.4 1.), namely, 36.46 
g. and 44 g., respectively. Hence the former gas is a little lighter. 

Remembering that the G.M.V. of air weighs 28.955 g. (Table, 
p. 101), we can compare the weight of any gas with that of air in 
the same way. What are the relative weights of acetylene (C 2 H 2 , 
p. 105) and sulphur dioxide (S0 2 ) as compared with air? The 
G.M.V. cube holds formula-weights of the first two, namely 26 g. 
and 64 g., and 28.955 g. of air. Hence acetylene is a little lighter 
than air, and sulphur dioxide more than twice as heavy (see 
Exercise 6). 

Exercises. — 1. What volume of oxygen at 10° and 750 mm. is 
obtainable by heating 50 g. of potassium chlorate (pp. 116, 150)? 

2. What volume of oxygen at 20° and 760 mm. is required to 
convert 16 g. of iron into dehydrated rust (Fe 2 3 ) (p. 150)? 

3. Write out the molecular equations for the interactions of 
methane and chlorine giving CH 3 C1; and for the burning of 
phosphorus (vapor) in oxygen (p. 105). Deduce the volume re- 
lations of the initial substances, and of the products, at various 
temperatures in each case. 

4. Write out the molecular equations for the interactions of 
acetylene and oxygen (p. 105), and of alcohol vapor (b.-p. 78°) and 
oxygen. Deduce the volume relations of the initial substances and 
of the products at 0° and at 100° in each case. 

5. The molecular weight of cyanogen is 52.08. What is its den- 
sity referred to air, and what the weight of 1 1. at 0° and 760 mm.? 
It contains 46.08 per cent carbon and 53.92 per cent nitrogen. 



CALCULATIONS 153 

What is the formula of the substance (p. 45)? Exploded with 
oxygen it forms carbon dioxide and free nitrogen. What will be 
the relative volumes of the materials before and after the inter- 
action (p. 151?) 

6. What are the relative weights of equal volumes of hydrogen 
sulphide (H 2 S), and hydrogen iodide (HI), compared with air 
(p. 152)? 



CHAPTER XII 

CHLORINE 

Chlorine was first recognized as a distinct substance by Scheele 
(1774). He obtained it from salt by means of manganese dioxide, 
using the method described below. It was supposed to be a com- 
pound containing oxygen until Davy (1809-1818) demonstrated 
that it was an element. 

Occurrence. — Chlorine does not occur free in nature. There 
are, however, many compounds of it to be found in the mineral 
kingdom. Sea-water contains a number of chlorides in solution. 
Of the 3.6 per cent of solid matter in sea-water, nearly 2.8% is 
sodium chloride NaCl. During past geological ages the evapora- 
tion of sea-water has led to the formation of immense deposits of 
the compounds usually found in such water. Thus, at Stassfurt, 
such strata attain a thickness of over a thousand feet. Certain 
layers of these strata are composed mainly of sodium chloride 
(rock salt). In other layers potassium chloride (sylvite), an in- 
dispensable fertilizer, and other compounds of chlorine, occur. 

Preparation. — Chlorine cannot be obtained with the same 
ease as oxygen. There are only a few chlorides, such as those of 
gold and platinum, which lose chlorine when heated, and they are 
too expensive or difficult to make for laboratory use. We employ 
therefore methods like those used for the preparation of hydrogen 
(cf. p. 53). We may (1) decompose any chloride by means of 
electricity, just as, to get hydrogen, we electrolyzed a dilute acid 
(p. 55). Or (2) we may take some inexpensive compound of 
chlorine, such as hydrogen chloride (HC1), and by means of some 
simple substance which is capable of uniting with the other con- 
stituent — here oxygen serves the purpose — secure the liberation 
of the element. Or (3) — and this turns out to be the most con- 
venient laboratory method — we may use a more complex action. 

154 



CHLORINE 



155 



Electrolysis of Chlorides. — Hydrogen chloride and those 
chlorides of metals which are soluble in water are all decomposed 
when a current of electricity is passed through the aqueous solu- 
tion. They yield chlorine at the positive electrode. The other 
constituent, the hydrogen (Fig. 65), manganese, or whatever it may 
be, is liberated at the negative 
wire. Since the chlorine is solu- 
ble in water, the effervescence 
due to its release is not notice- 
able until the liquid round the 
electrode has become saturated 
with the gas : Cl 2 (dslvd) ^ Cl 2 
(gas) . The shape of the appa- 
ratus keeps the two products 
from mingling. The presence 
of the chlorine in the liquid at 
the positive end may be shown 
by a suitable test (p. 161). 

In commerce chlorine is now 
obtained chiefly by this method, 
sodium chloride or potassium 
chloride being the source of the element. Electrodes of artificial 
graphite are used, as most other conductors unite with the 
chlorine. The potassium or sodium, as the case may be, 
travels towards the negative electrode, but is not liberated. 
Instead, potassium or sodium hydroxide (q.v.) accumulates in 
the solution round the plate and hydrogen escapes. The chlo- 
rine is released at the positive electrode, as usual. The hydro- 
gen, the hydroxide and the chlorine all find commercial applica- 
tions. The chlorine is either liquefied by compression in soeel 
cylinders or is employed at once for making bleaching powder 
(see index). 




Fig. 65. 



Action of Free Oxygen on Chlorides. — Sodium chloride is 
the cheapest source of chlorine,[but oxygen does not interact with 
it even at a high temperature. By treating the sodium chloride 
with sulphuric acid, therefore, the chlorine is first transferred into 
combination with the hydrogen of the acid, giving hydrogen 
chloride (p. 141). In order to liberate chlorine from the hydrogen 



156 COLLEGE CHEMISTRY 

chloride, we may then combine the hydrogen with oxygen obtained 
from the air. 

Skeleton: HC1 + <± H 2 + CI. 

Balanced: 2HC1 + <=± H 2 + 2C1. 

Molecular: 4HC1 + 2 <=± 2H 2 + 2C1 2 . 

The two gases interact so slowly, however, that a contact agent 
must be employed. The mixture of air and hydrogen chloride is 
passed over pieces of heated pumice-stone (Fig. 66) or broken 
brick previously saturated with cupric chloride solution. A tem- 
perature of about 370° is used. 
*%= m ^^^^^^s^.^ 4gp ==» Furthermore, the action is re- 
versible (read the equation 
backwards) and equilibrium is 
reached when 80 per cent of the hydrogen chloride has been 
decomposed. Hence 20 per cent of this gas passes on unchanged. 
Only 80 per cent of the hydrogen chloride and oxygen are changed 
into steam and chlorine, because the latter substances are continu- 
ously interacting to reproduce hydrogen chloride and oxygen. If 
one substance could be separated (p. 143) from the other, to pre- 
vent the backward action, the yield would be raised to 100 per 
cent. In the product, the chlorine is mixed with steam and with a 
very large volume of nitrogen which entered with the oxygen, as 
well as with unused hydrogen chloride, so that, for making the 
pure substance, this method (Deacon's process) is quite unsuitable. 
Bleaching powder, however, can be made by its means. 

The relative volumes in this reaction (see p. 150) are indicated 
by the numbers of molecules in the equation. Four volumes of 
hydrogen chloride and one volume of oxygen give two volumes of 
steam and two volumes of chlorine. 

The above action is spoken of as an oxidation. It is true that 
no oxygen is actually introduced into the hydrogen chloride as 
a whole. The removal of hydrogen from combination with the 
chlorine is, however, the first step towards the introduction of 
oxygen into combination with the latter, and is essentially an 
oxidation. 

Action of Combined Oxygen upon Chlorides. — The best 
laboratory method for making chlorine is to place some solid 



CHLORINE 



157 




Fig. 67. 



potassium permanganate in a flask, arranged like that in Fig. 67. 
Concentrated hydrochloric acid (an*aqueous solution of hydrogen 
chloride), diluted with one- 
third of its volume of water, 
is allowed to fall upon the com- 
pound drop by drop from the 
dropping funnel. The action 
is very rapid, the acid is ex- 
hausted almost as fast as it 
falls, and so the stream of gas 
can be stopped by simply clos- 
ing the stopcock. The gas 
is passed through a washing 
bottle containing water, in 
order to remove any hydrogen 
chloride which may be carried 
over. It may be dried, if 
necessary, in a second washing bottle containing concentrated 
sulphuric acid. It cannot be collected over water on account of 
its solubility, so that jars are usually filled with it by upward 
displacement of air. 

Skeleton: KMn0 4 + HC1 -» H 2 + KC1 + MnCl 2 + CI. 

The O4, being all converted into water, requires 8H, and therefore 
8HC1, for the action. The two metals, potassium and manganese, 
give their respective chlorides, KC1 and MnCl 2 . This uses 3C1, and 
hence 5C1 remains over to be liberated: 

Balanced: KMn0 4 + 8HC1 -> 4H 2 + KC1 + MnCl 2 + 5C1. 
Molecular: 2KMn0 4 + 16HC1 -> 8H 2 + 2KC1 + 2MnCl 2 + 5C1 2 . 

The combined oxygen of the permanganate has oxidized the hydro- 
gen chloride, just as did the free oxygen in Deacon's process. 

Other Means of Oxidizing Hydrogen Chloride, — Many 
other compounds of oxygen with metals interact with hydro- 
chloric acid to give free chlorine. Lead dioxide Pb0 2 , potassium 
chlorate KCIO3, potassium dichromate K 2 Cr 2 07, and manganese 
dioxide Mn0 2 , are of this nature. The last, being inexpensive, is 
commonly used in making chlorine. Being an insoluble substance, 



158 COLLEGE CHEMISTRY 

however, the manganese dioxide acts much more slowly than does 
the potassium permanganate, which is soluble. A large amount of 
the materials, and the aid of heat, are required to secure a rapid 
stream of chlorine. 

Manganese Dioxide and Hydrogen Chloride. — The action 
of manganese dioxide upon hydrochloric acid is an instructive one. 
It is a general rule, of which we shall meet many applications, that 
when an acid interacts with an oxide of a metal, there are two con- 
stant features in the result, namely: (1) The oxygen of the oxide 
combines with the hydrogen of the acid to form water, and (2) the 
metal of the oxide combines with the acid radical of the acid accord- 
ing to the valences of each. Here the skeleton equation should be 
Mn0 2 + HC1 -> H 2 + MnCU. With 2 , to form water, 4HC1 is 
required, and the product is 2H 2 0. Hence the equation is 

Balanced: Mn0 2 + 4HC1 -* 2H 2 + MnCU- 

This is what happens in the first place. The products actually 
obtained, however, are water, manganous chloride MnCl 2 and 
chlorine. The manganese tetrachloride can be preserved by cool- 
ing the mixture. It is decomposed by the heating, the chlorine 
escapes, and the other two products remain in the vessel. 

Mn0 2 + 4HC1 -» 2H 2 + MnCl 2 + Cl 2 . (1) 

We owe the chlorine to the fact that the tetrachloride is unstable. 
If we had used manganous oxide MnO, we should have had a 
double decomposition: 

MnO + 2HC1 -> H 2 + MnCl 2 , (2) 

but we should have got no chlorine. Perhaps the simplest way to 
describe the difference between these two actions is in terms of the 
valence of the manganese. In Mn IV 02 n the element is quadriva- 
lent. This means that its atomic weight professes to be able to 
hold four atomic weights of a univalent element. The four valences 
of oxygen (20 n ) can do the same thing. In equation (1) the 
oxygen fulfils this promise by taking 4H 1 . But the Mn 17 can hold 
only 2C1 1 , permanently, and lets the other 2C1 1 go free. In other 
words, the valence of the atomic weight of manganese changes in the 
course of the action. In equation (2), on the other hand, the 
manganese is bivalent to start with (Mn n O I]C ), and is able to retain 



CHLORINE 159 

the amount of chlorine (2C1 1 ) equivalent to O u . Actions like that 
of manganese dioxide in (1) are classed as oxidations. The hydro- 
gen chloride, or rather half of it, is oxidized. A graphic mode of 
writing may make this remark clearer : 

z + 2HC1 -4 H 2 + Mn u Cl 2 
Mn Iv f 

^0 + 2HCl->H 2 + Cl 2 

The upper half is a double decomposition, the lower an oxidation 
by half the combined oxygen of the dioxide. The same explana- 
tion applies to the interaction of lead dioxide with hydrochloric 
acid. 

Physical Properties. — Chlorine differs from the gases we 
have encountered so far in having a strong greenish-yellow tint, 
a fact which gave rise to its name (Gk., pale green), and having a 
powerful, irritating effect upon the membranes of the nose and 
throat. 

Density (H = 1), 35.79. Boiling-point (liq.), -33.6°. 

Weight of 22.4 1., 72.13 g. Melting-point (solid), -102°. 

Sol'ty in Aq (20°), 215 vols, in 100. Vap. tension (liq.) 0°, 3.66 atmos. 

Crit. temp., +146°. Vap. tension (liq.) 20°, 6.62 atmos. 

Since the G.M.V. of air weighs 28.95 g., chlorine is two and a half 
times heavier. In solubility it stands between slightly soluble 
gases, like oxygen and hydrogen, and those which are extremely 
soluble. It can be collected over hot water or a strong solution of 
salt. 

Chlorine was first liquefied by Northmore (1806). It forms a 
yellow liquid which, contained in steel cylinders lined with lead, 
is now an article of commerce. On being cooled below — 102°, it 
gives a pale-yellow solid. 

In recalling the physical properties of a gas, remember that six 
(p. 31) are required: color, taste, odor, density, solubility, lique- 
fiabihty. 

Chemical Properties. — Chlorine is at least as active a sub- 
stance as is oxygen. It presents a more varied array of chemical 
properties than does that element. The binary compounds are 
called chlorides. 



160 COLLEGE CHEMISTRY 

Combines with Metals. — Powdered antimony (cold), when 
thrown into chlorine, unites with it to form the chloride SDCI3, 
which appears partly as vapor and partly as glowing particles. 
Balanced: Sb + 3C1 -> SbCl 3 . 

Molecular: 2Sb + 3C1 2 -> 2SbCl 3 . 

Copper, in the condition of thin leaf commonly used for gilding 
(Dutch-metal), catches fire when thrust into the gas, giving a fog 
of solid cupric chloride CuCl 2 . Sodium burns brilliantly, giving 
a cloud of sodium chloride. The union of a metal like sodium 
and a colored, irritating gas to give a mild household article, like 
common salt, illustrates the extraordinary nature of chemical 
change. All the familiar metals, with the exceptions of gold and 
platinum, combine readily with chlorine. 

When metals (like copper and iron) and chlorine are first thor- 
oughly freed from moisture, combination no longer occurs. A 
trace of water is required in these, as it is in many other chemical 
actions, as a contact agent. Hence, the chlorine, before being 
compressed into steel cylinders, must be freed entirely from water 
vapor (see Detinning). 

Combines with Hydrogen. — A jet of hydrogen burns 
vigorously in chlorine, producing hydrogen chloride HC1. The 
union of the gases, when a mixture of them is kept cold and in the 
dark, is too slow to be perceived. On exposure to diffused light, 
however, they unite slowly, while a sudden flash of sunlight or the 
burning of a magnesium ribbon causes instant explosion. The 
effect of the light is catalytic. 

Interacts with Compounds Containing Hydrogen. — 

When a lighted taper is plunged into chlorine it continues to burn, 
but a dense cloud of soot (free carbon) rises from the flame. Blow- 
ing the breath into the jar then gives the fog which shows the 
presence of hydrogen chloride. Thus the presence of hydrogen 
and carbon in the wax is proved. We learn, also, that chlorine 
has little tendency to combine with carbon, for this element goes 
free. A few drops of warm turpentine, poured upon a strip of 
paper, when placed in chlorine give a violent reaction and a cloud 
of finely divided carbon bursts forth. 

C 10 H 16 + 8C1 2 -> 16HC1 + IOC, 



CHLORINE 161 

Elements Displaced by Chlorine. — The action on turpen- 
tine is a displacement of the carbon by the chlorine. Of the same 
nature is the action of chlorine upon potassium iodide KI, dry or 
in solution. 

2KI + C1 2 ->2KC1 + I 2 . 

The iodine, when moist, is deep brown in color. A mere trace of 
chlorine, liberating a trace of iodine, gives no visible effect. But 
if some starch is present, even a trace of free iodine yields a deep 
blue color. This reaction is used as a test for chlorine, for free 
iodine from any source, and for starch (p. 3). To test for chlorine, 
strips of filter paper, dipped in starch emulsion (starch boiled with 
much water and cooled) to which a few drops of potassium iodide 
have been added, are used. Combined iodine, as in potassium 
iodide, has no effect upon starch. Combined chlorine, as in 
sodium chloride, has no action upon the prepared strips of paper — 
free chlorine is required. 

Action Upon Water. — We have seen that chlorine seizes the 
hydrogen in turpentine. We have also learned that it combines 
with the hydrogen in steam, reversing Deacon's process to the 
extent of 20 per cent. It also acts upon cold water, when dissolved 
in the latter, although in a similarly incomplete way. The sub- 
stances formed are hydrochloric acid and hypochlorous acid HCIO : 

H 2 + Cl 2 +± HC1 + HCIO. 

With half -saturated chlorine-water at 10° — that is, water con- 
taining an equal volume of chlorine gas — 33 per cent of the 
chlorine is changed into the acids. Thus, chlorine-water (the 
solution) is a mixture containing dissolved chlorine and two acids. 
Hypochlorous acid (q.v.) is of especial interest because it is a very 
active substance, with powerful oxidizing qualities, and bleaches 
dyes by decomposing them. 

The action comes to a standstill when one-third completed, 
because the two acids interact to reproduce chlorine and water 
(read the equation backwards). The action is reversible. When 
the solution is exposed to sunlight, the hypochlorous acid decom- 
poses and oxygen gas is liberated and escapes: 

2HC10 -> 2HC1 + 2 T • 



162 COLLEGE CHEMISTEY 

Since this removes the hypochlorous acid, on whose interaction 
with the hydrogen chloride the reverse action depends, the for- 
ward action proceeds under continuous illumination gradually to 
completion. Hence the aqueous solution of chlorine must be kept 
in the dark, since otherwise, after a time, a dilute solution of 
hydrogen chloride alone remains. 

The reader should note here the displacement of the equilibrium, 
a chemical one in this case, in consequence of the annulment of one 
of the opposing tendencies (p. 90). Through the destruction of 
the hypochlorous acid, one of the tendencies, namely that repre- 
sented in the backward action, becomes inoperative. The for- 
ward action is not itself assisted, but it is no longer impeded, and 
so proceeds to completion. 

Action by Substitution. — When actions like that on tur- 
pentine — that is on compounds containing carbon and hydrogen 
— are moderated by altering the conditions, the decomposition 
is not so complete. Using a lower temperature is effective. Thus, 
if methane CH4 (marsh-gas), the chief component of natural gas, 
is mixed with chlorine and exposed to sunlight, a slower action 
occurs, of which the first stage consists in the removal of one unit 
weight of hydrogen and the substitution of chlorine for it according 
to the following equation: 

CH4 + Cl 2 -> CH 3 C1 + HC1. 

The process may continue further by the substitution * of chlorine 
for the units of hydrogen one by one until carbon tetrachloride 
CCU is finally formed. 

The action on water is a substitution. 

Combines with Non-metals. — Phosphorus burns in chlorine 
with a rather feeble light, producing primarily phosphorus tri- 

* Substitution resembles displacement (p. 55) in that an element and a 
compound interact, and the element takes the place of one unit in the com- 
position of the latter. In the above action, one unit of chlorine takes the 
place of one unit of hydrogen. But the latter is not liberated; it combines 
with another unit of chlorine. The action resembles double decomposition, 
excepting that one of the substances is not a compound, but a diatomic ele- 
ment. The name used is intended to fix the attention on the compound and 
on the fact that one unit has been substituted for another in it. This concep- 
tion is a favorite one in the chemistry of compounds of carbon. 



CHLORINE 163 

chloride PC1 3 , a liquid (b.-p. 74°). If excess of chlorine is present, 
then, as the trichloride cools, it combines to form the solid penta- 
chloride PCI3. Sulphur, when heated, unites more slowly, giving 
sulphur monochloride S 2 C1 2 , a liquid used in vulcanizing rubber. 
Chlorine does not combine directly with carbon, nitrogen, or oxy- 
gen, although compounds with those elements can be made in- 
directly. With the helium group of elements (q.v.), it forms no 
compounds. 

Combines ivith Compounds. — Chlorine unites with many 
compounds. Thus, one of the oxides of carbon, carbon monoxide 
CO, when mixed with chlorine and exposed to sunlight gives drops 
of a volatile liquid (b.-p. 8.2°) known as phosgene COCl 2 . 

When chlorine-water is cooled with ice, a compound, chlorine 
hydrate C1 2 ,8H 2 crystallizes out. Faraday (1823) placed this 
substance in the closed limb of a A-tube, 
sealed the open end, and placed the empty 
limb in cold water (Fig. 68). When the 
hydrate was gently warmed, chlorine gas 
was given off and was liquefied by its own 
pressure in the cold part of the tube. 

Fig. 68. 

Chemical Relations of the Element.* — In the chlorides, 
an atomic weight of chlorine is equivalent to one atomic weight 
of hydrogen or of sodium. The element is, therefore, univalent 
(p. 62). It never shows any higher valence than this, save in 
its oxygen compounds (see Chap. XXIII). The oxides of chlo- 
rine interact with water to give acids, and the element is, there- 
fore, to be classed as a non-metal (p. 94). It belongs to that 
group of the non-metals called the halogens, as a consideration of 
some others of its relations will show (see Chap. XV). 

* In accordance with the distinction that must be drawn (p. 16) between 
the element as a variety of matter in combination, and the elementary sub- 
stance or free form of the element, and to avoid a common source of con- 
fusion, we shall always give only the behavior of the elementary substance 
under the title chemical properties. The characteristics which distinguish 
the compounds of the element, as a class, from, or relate them as a class to the 
compounds of other elements will then appear in a separate section under 
the title "Chemical relations" (see pp. 192, 208). 




164 



COLLEGE CHEMISTRY 



Uses of Chlorine. — Large quantities of chlorine are manu- 
factured for the preparation of bleaching materials and disinfect- 
ing agents. In disinfection, the minute germs of disease and 
putrefaction are acted upon either by the chlorine or by the hypo- 
chlorous acid formed by its interaction with water, and instantly 
their life is destroyed. 



Chlorides. — The chlorides are described individually under 
the other element which each contains. The majority of the 
chlorides of the metals are easily soluble in water. The chief 
exceptions are silver chloride AgCl, mercurous chloride (calomel) 

HgCl, cuprous chloride CuCl, and lead 
chloride PbCl 2 . The last of these is on 
the border line as regards solubility. An 
appreciable amount dissolves in cold 
water, and a considerable amount in 
boiling water (see Table of Solubilities, 
inside the cover at the front of this 
book). For the various modes of pre- 
paring chlorides see p. 146. 




Composition of Hydrogen Chlo- 
ride. — Being now familiar with both 
hydrogen and chlorine, we may take up 
the question of the proportion by vol- 
ume in which the constituents unite, 
and the relation of this to the volume 
of the resulting hydrogen chloride. 

The decomposition of the solution of 
hydrogen chloride in water by means of 
the electric current proves that the gases 
are liberated in equal volumes. Brown- 
lee's apparatus for demonstrating this is shown in Fig. 69. The cen- 
tral part is the same as in Fig. 27, but, when the three-way stop- 
cock is closed, the gases go to right and left, and displace the 
liquid in two outside tubes. The equal rate at which this takes place 
on both sides proves that the gases are generated in equal volumes. 
In order to ascertain the relation between the volumes of the 
constituents and that of the product, we may unite the gases and find 



Fig. 69. 



CHLORINE 



165 



out whether any change in volume occurs. A tube with thick walls 
(Fig. 70) is filled with the mixed gases obtained by electrolysis. 
By dipping one end of the tube under mercury and . 

opening the lower stopcock, it is seen that no gas leaves n? 
and no mercury enters. After the mixture has been f^ 
exploded, by the light from burning magnesium, the 
same test is repeated with the same result. The pres- 
sure has therefore remained equal to that of the at- 
mosphere. Hence there has been no change in volume 
as the result of the union. It appears, therefore, that 



1 vol. hydrogen + 1 vol. chlorine -* 
2 vols, hydrogen chloride, 

a result in harmony with Gay-Lussac's law (p. 98). 



Fig. 70. 



Confirmation of the Formulae Cl 2 and H 2 . — According to 
Avogadro's law, there are equal numbers of molecules in equal 
volumes of these gases. When hydrogen and chlorine combine, 
one volume of each of these gases gives two volumes of hydrogen 
chloride. Let us imagine the experiment to be made with minute 
volumes holding one hundred molecules each: 

Hydrogen Chloride Hydrogen Chlorine 

came from 



100 


100 



100 


+ 


100 



The 200 molecules of hydrogen chloride must contain at least 200 
fragments of chlorine, since there is a sample in each molecule. 
Now the 200 fragments of chlorine came from a volume contain- 
ing only 100 molecules of chlorine. Each of these must therefore 
have been split in the chemical action. The same is true of each 
molecule of hydrogen. Hence the molecules of free hydrogen 
and free chlorine contain at least two atoms. If we consider the 
molecular formula of a substance as representing one molecule, 
the equation for this action is: 

H 2 + Cl 2 -> 2HC1. 

There are two molecules on each side of the equation, and this 
corresponds with the fact that there is no change in the total 
volume. 



166 COLLEGE CHEMISTRY 

Classification of Chemical Interactions and Exercises 
Thereon. — So far we have defined ten more or less distinct kinds 
of chemical change, seven differing in mechanism: Combination 
(p. 7), decomposition (p. 14), dissociation (p. 93), displacement (p. 
55), substitution (p. 162), double decomposition (pp. 142, 147), and 
internal rearrangement (p. 148); and three others: oxidation (pp. 36, 
156, 158), reduction (pp. 37, 59), and electrolysis (pp. 55, 155). 
Illustrations of all but one of these will be found in the present 
chapter. Some actions belong to one of the first seven, and also 
to one of the three other classes. The ability readily to classify 
each phenomenon, as it comes up, requires precisely that grasp 
of the framework of the science which the reader must seek speedily 
to attain. For example, let him classify the following actions: 1. 
action of potassium on water; 2. of heat on potassium chlorate; 
3. of chlorine on metals; 4. of chlorine on turpentine; 5. of 
chlorine on potassium iodide; 6. of chlorine on methane; 7. of 
carbon monoxide and chlorine; 8. of sunlight on hypochlorous 
acid; 9. of sulphuric acid on salt; 10. of zinc oxide and hydro- 
chloric acid; 11. of zinc on hydrochloric acid. 

12. In the interactions of potassium permanganate and of man- 
ganese dioxide, respectively, with hydrochloric acid, what fractions 
of the whole chlorine are liberated? What are the commercial 
advantages of the use of salt and sulphuric acid with the manganese 
dioxide? 

13. In view of the explanations given, define the general nature 
of the substances (p. 157) which may be used to oxidize hydro- 
chloric acid. 

14. What are the relative volumes of the gaseous interacting 
substances and products in the following reactions : (a) turpentine 
vapor and chlorine; (6) methane and chlorine; (c) phosphorus 
vapor and chlorine; (d) carbon monoxide and chlorine. 



CHAPTER XIII 

ENERGY AND CHEMICAL CHANGE 

In describing chemical changes, the fact that heat was evolved 
has frequently been mentioned. In several instances]a current of 
electricity has been used to produce chemical change. It is now 
necessary to collect these scattered facts and classify them for 
future use. 

Physical Accompaniments of Chemical Change, — When 
iron and sulphur combined (p. 13), and when iron burned in oxy- 
gen or copper in chlorine, much heat was developed. On the 
other hand, the decomposition of mercuric oxide, as was pointed 
out (p. 14), owed its continuance to the persistent application of 
heat and ceased as soon as the source of heat was withdrawn. 
Here, apparently, heat was consumed during the progress of the 
change, and the chemical action was limited by the amount of 
heat supplied. The production or consumption of heat may, there- 
fore, be a feature of chemical change. 

In the burning of iron or magnesium in oxygen, and in the 
actions of chlorine on copper and turpentine, light was also pro- 
duced. Conversely, silver chloride (p. 147) can be kept any length 
of time in the dark, but in sunlight it becomes first bluish and 
then brown, simultaneously giving off chlorine gas and finally 
leaving only silver as a fine powder. Silver bromide or iodide, in 
photographic plates, films, and paper, is changed by light in a 
similar way, liberating the bromine or iodine. It would appear, 
therefore, that light may be given out or consumed in connection 
with chemical change. 

We have seen (p. 155) that a current of electricity may be em- 
ployed to decompose hydrochloric acid and other chlorides, and 
the battery, or other source of the current must be kept going or 
the chemical change stops. The inverse of this is likewise familiar. 
If we place in dilute sulphuric acid a stick of the metal zinc, we 

167 



168 



COLLEGE CHEMISTRY 



i\\'^r/^^ 



find that hydrogen is given off (Fig. 71), that the zinc goes into 
solution as zinc sulphate (p. 53), and that a large amount of heat 
is developed. If zinc in fine particles, with 
much surface, is used, the liquid may even 
rise spontaneously to the boiling-point. 
This form of the action produces heat. If, 
however, we attach the same stick of zinc 
to a copper wire and, having provided a 
plate of platinum also connected with a wire, 
immerse the two simultaneously in the acid 
(Fig. 72), then a galvanometer, with which 
the wires are connected, shows at once the 
passage of a current of electricity round 
the circuit. Exactly the same chemical 
change goes on as before. The sole differ- 
ence is that the gas appears to arise from 
the surface of the platinum. It is easy to 
show, however, that the platinum by itself 
is not acted upon by dilute acids and, in 
this case, undergoes no change whatever; it serves simply as a 
suitable conductor for the electricity. Here, then, in place of the 




Fig. 71. 




Fig. 72. 



heat which the first plan produced, we get an electric current. 
The arrangement is, in fact, a battery-cell, for a battery is a 
system in which a chemical action which would otherwise give 



ENERGY AND CHEMICAL CHANGE 169 

heat furnishes electricity instead. Thus, electrical energy may be 
consumed or produced in connection with a change in composition. 

Even violent rubbing in a mortar, in the case of some substances, 
can effect an appreciable amount of decomposition in a few min- 
utes. In this way silver chloride can be separated into silver and 
chlorine, just as by light. It is the mechanical energy which is the 
agent, and part of it is consumed in producing the change, and 
only the balance appears as heat. Conversely, the production of 
mechanical energy, as the result of chemical change, is seen in 
the behavior of explosives and in the working of our muscles. 
Thus, mechanical energy may be used up or produced in chemical 
changes. 

Summing up our experience, we may state that no change in 
composition occurs without some accompaniments, such as the 
production or consumption of heat, light, electrical energy, or, in 
some cases, mechanical energy. 

Classification of the Accompaniments of Change in Com- 
position: Energy. — The problem of classifying (i.e., placing in a 
suitable category) things like heat, light, and electricity has occu- 
pied much attention. In all changes in composition, one of these 
natural accompaniments is given out or absorbed, sometimes in 
great amount, yet in none is any alteration in weight observed.* 
There are many things which are real, however, even if they are 
not affected by gravitation. In the present instance we reason 
as follows : 

A brick in motion is different from a brick at rest. The former 
can do some things that the latter cannot. Furthermore, we can 
easily make a distinction in our minds. The brick can be deprived 
of the motion and be endowed with it again. Thus, we can get 
the idea of motion as a separate conception. Similarly, we observe 
that a piece of iron behaves differently when hot, and when cold, 
when bearing a current of electricity, and when bearing none. 
We conceive then of the brick or the iron as having a certain 
amount and kind of matter which is unalterable, and as having 
motion, heat, or electricity added to this or removed. Thus, we 
describe our observations by using two categories, one of which 

* Electrons (q.v.) do possess mass, but it is very small compared with that 
of the materials concerned. 



170 



COLLEGE CHEMISTEY 



includes the various kinds of matter, and the other, various things 
whose association with matter seems to be invariable and is often 
so conspicuous. The latter we call the forms of energy. 

The Practical Importance of Energy in Chemistry, — 

The absorption or liberation of energy accompanying a chemical 
transformation of matter is often, of the two, the more important 
feature. We do not burn coal in order to manufacture carbon 
dioxide gas. We are glad to get rid of the material product through 
the chimney. It is the heat we want. We do not buy gasoline 
(petrol) for an automobile in order to obtain various gases to ex- 
pel through the muffler. We really pay for the mechanical energy. 
It is the same with burning illuminating-gas or magnesium powder 
when we want light, and with eating food, which we do, chiefly, to 
get energy to sustain our activity. We do not run electricity for 
hours into a storage battery in order to make a particular compound 

(lead dioxide, for example), but in order 
to save and store the energy for future 
use. In industry and life fully half the 
total amount of chemical change in- 
volved is set in motion by us, solely on 
account of the .energy changes it in- 
volves. But the production of energy 
in chemical change is not only thus of 
practical importance; it is also of scien- 
tific interest, as will be seen in the sec- 
tion on energy and chemical activity 
(see below). 




Fig. 73. 



Interconvertibility of Forms of 
Energy: Conservation. — At first 
sight, these accompaniments of matter 
seem to be quite unrelated. But a relation between them can be 
found. If the heat of a Bunsen flame or of the sun is brought 
under a hot-air motor (Fig. 73) violent motion results. Again, if 
the motor is connected with a dynamo, electricity may be generated. 
Still again, if the current from the dynamo flows through an incan- 
descent lamp, heat and light are evolved, Conversely, when motion 
of the hot-air motor is impeded by a brake, heat appears. When 



ENERGY AND CHEMICAL CHANGE 171 

a current of electricity is run through the dynamo, the armature 
of the latter turns and motion results. But the most significant 
facts are still to be mentioned. The heat absorbed by the motor 
is found to be greater when the machine is permitted to move and 
do work, than when it is not. Thus, it is found that when work is 
done some heat disappears, and this heat is, in fact, transformed 
into work. Similarly, when the poles of the dynamo are properly 
connected and electricity is being produced, and only then, motion 
is used up. This is shown by the effort required to turn the arma- 
ture under these circumstances, and the ease with which it is 
turned when the circuit is open. So, with a conductor like the 
filament in the lamp, unless it offers resistance to the current and 
destroys a sufficient amount of electrical energy, it gives out neither 
light nor heat. Finally, motion gives no heat unless the brake is 
set, and effort is then demanded to maintain the motion. These 
experiences lead us to believe that we have here a set of things 
which are fundamentally of the same kind, for each form can be 
made from any of the others. We have, therefore, invented the 
conception of a single thing, of which heat, fight, electricity, and 
motion are forms, and to it we give the name energy: energy is 
work and every other thing which can arise from work and be con- 
verted into work (Ostwald) . 

Closer study shows that equal amounts of electrical or mechani- 
cal energy always produce equal amounts of heat. No loss is 
ever observed in the transformations of energy, any more than 
in the transformations of matter. Hence, J. R. Mayer (1842), 
Colding (1843), and Helmholtz (1847) were led independently to 
the conclusion that in a limited system no gain or loss of energy 
is ever observed. This brief statement of the results of many ex- 
periments is called the law of the conservation of energy. 

Application of the Conception of Energy in Chemistry. — 

At first sight it looks as if the statement that energy is conserved 
is not applicable in chemistry. Heat and electricity, for example, 
seem to be produced and consumed, in connection with changes in 
composition, in a mysterious manner. We trace light in an in- 
candescent lamp back to the electricity, and this in turn to the 
mechanical energy, and this again to the heat in the engine. But 
what form of energy gave the heat developed by the combustion 



172 COLLEGE CHEMISTRY 

of the coal under the boiler, or by the union of iron and sulphur 
in our Jirst experiment? Since we do not perceive any electricity, 
light, heat, or motion, in the original materials, and yet wish to create 
an harmonious system, we are bound to conceive of the iron and the 
sulphur, and the coal and the air, as containing another form of 
energy, which we call internal energy. Similarly, when heat is used 
up in decomposing mercuric oxide, or light in decomposing silver 
chloride, we regard the energy as passing into, and being stored 
in the products of decomposition in the form of internal energy. 

These conclusions compel us, for the sake of consistency, to 
think of all our materials as repositories of energy as well as of 
matter, each of these two constituents being equally real and 
equally important. A piece of the substance known as "iron" 
must thus be held to contain so much iron matter and so much 
internal energy. So ferrous sulphide contains sulphur matter, 
iron matter, and internal energy. Thus, by a substance we mean 
a distinct species of matter, simple or compound, with its appro- 
priate proportion of internal energy. 

In the course of this discussion it has become clear that it is 
characteristic of chemical phenomena that, besides a change in the 
nature of the matter, there is always an alteration in the amount of 
internal energy in the system. This alteration involves the produc- 
tion of internal energy from, or the transformation of internal 
energy into some other form of energy. 

Energy and Chemical Activity. — Other things being equal, 
actions in which there is a relatively large loss of internal energy 
proceed rapidly; that is to say, in them a large proportion of the 
material is changed in the unit of time. Those in which less en- 
ergy is transformed proceed more slowly. The speed of the chemi- 
cal change, and the quantity of energy available because of it, are 
closely related. Now, we are accustomed to speak of materials 
which, like iron and sulphur, interact rapidly and with liberation 
of much energy as "chemically active." Thus, relative chemical 
activity may be estimated, (1) by observing the speed of a change 
(see below), or, in many cases (2) by measuring the heat developed 
(see Thermo-chemistry, below), or (3) by ascertaining the electro- 
motive force of the current, when the materials are arranged in 
the form of a battery-cell (see Chap. XXXIX). 



ENERGY AND CHEMICAL CHANGE 173 

It is evident that the chemical activity of a given substance will 
not be the same towards all others. Thus, iron unites much more 
vigorously with chlorine than with sulphur and, with identical 
amounts of iron, more heat is liberated in the former case than in 
the latter. With silver, sodium, and many other substances, iron 
does not unite at all. One of the tasks of the chemist is to make 
such comparisons as this. He calls the results the specific chemical 
properties of the substances in question. 

The " Cause " of Chemical Activity or Affinity. — The 

reader will undoubtedly be inclined to inquire whether we can 
assign any cause for the tendency which substances have to undergo 
chemical change. Why do iron and sulphur unite to form ferrous 
sulphide, while other pairs of elements taken at random will fre- 
quently be found to have no effect upon one another under any 
circumstances? The answer is that we do not know. Questions 
like this have to go without answer in all sciences. What is the 
cause of gravitation? We know the facts which are associated 
with the word — the fact that bodies fall towards the earth, for 
example — but why they fall we are unable" to say. So, with 
chemical change, we can state all the facts we know about it, but 
even then we cannot say why it takes place. 

The word cause was employed in the heading of this section, and 
it will be observed that no cause was found. This is the invariable 
rule in physical or chemical phenomena. We know of no causes, 
in the sense in which the word is commonly employed. 

The word cause has only one definite use in science. When we 
find that thorough incorporation of the three materials is needed to 
secure good gunpowder, we say that the intimate mixing is a cause 
of its being highly explosive. By this we simply mean that intimate 
mixture is a necessary antecedent of the result. A cause is a condition 
or occurrence which always precedes another condition or occurrence. 

The Speed of Chemical Actions: a Means of Measuring 
Activity. — One means of measuring the relative chemical activi- 
ties of several substances is to observe the speed with which they 
undergo the same chemical change. Thus we may compare the 
activities of the various metals by allowing them separately to 
interact with hydrochloric acid and collecting and measuring the 



174 COLLEGE CHEMISTRY 

hydrogen liberated per minute by each. It will be seen, even in the 
roughest experiment, that magnesium is thus much more active 
than zinc. The comparison must be made with such precautions, 
however, as will make it certain that the conditions under which 
the several metals act are all alike. Thus, in spite of the heat 
evolved by the action, means must be used, by suitable cooling, to 
keep the temperature at some fixed point during the experiment, 
for all actions become more rapid when the temperature rises 
(p. 20). Again, the pieces of the various metals must be arranged 
so that equal surfaces are exposed to the acid in each case. It is 
found that the order in which this comparison places the metals is 
much the same as that in which they are placed by a study of 
other similar actions. A single table, showing the order of activ- 
ity (p. 60), suffices, therefore, for all purposes. 

Thermochemistry. — Chemical changes in which heat is 
liberated are called exothermal. Those in which heat is continu- 
ously absorbed (pp. 14, 167, 113) are called endothermal changes. 
Since the activities, or affinities of two substances (say, two metals) 
may often be measured by observing the amounts of heat liberated 
when each combines with a third substance (say, oxygen), it will 
be instructive now to consider some of the elementary facts of 
thermochemistry. 

The chemical interactions to be studied thermally are arranged 
so that they may be carried out in a small vessel which can be 
placed inside another containing water. The whole apparatus is 
called a calorimeter (Gk., heat-measurer). The heat developed 
raises the temperature of this water. Where gases like oxygen 
are concerned, a closed bulb of platinum forms the inner vessel. 
The quantity of heat capable of raising one gram of water one 
degree in temperature at 15° Centigrade is called a calorie. So that 
250 grams of water raised 1° would represent 250 calories, and 20 
grams of water raised 5° would represent 100 calories. 

Thermochemical Equations. — While in physics the unit 
of quantity is the gram, in chemistry the unit which we select is 
naturally a gram-atomic weight or a gram-molecular weight of the 
substance. Thus, the heat of combustion of carbon means the 
heat produced by combining twelve grams of carbon with thirty- 



ENERGY AND CHEMICAL CHANGE 175 

two grams of oxygen, and is sufficient to raise nearly 100,000 
grams of water one degree. This is expressed as follows: 
C + 2 ^C0 2 + 96,820 cal. 

In other words, the combustion of less than half an ounce of carbon 
will raise over two pounds (one kilogram) of water from 0° to the 
boiling-point. 

When the action is one which absorbs heat, this fact is indicated 
by the negative sign preceding the number of calories. Thus, the 
dissociation of 36 g. of water vapor into hydrogen and oxygen 
absorbs 28,800 cal. per gram of hydrogen liberated: 

2H 2 -> 2H 2 + 2 - 115,200 cal. 

If the action is reversible, as this one is, the heat absorbed when 
it proceeds in one direction is equal to that liberated when it goes 
in the other direction: 

2H 2 + 2 -» 2H 2 + 115,200 cal. 

An action which absorbs heat can take place only when heat or 
some other form of energy is furnished. Thus, the electrolysis of 
aqueous hydrochloric acid (p. 155) consumes electrical energy, 
which is equivalent in amount to the heat given out when hydrogen 
and chlorine unite to form hydrogen chloride, plus the heat liber- 
ated when the latter dissolves: 

H 2 + Cl 2 + Aq -» 2HC1, Aq + 78,800 cal. 

Answers to Possible Questions. — It is always found that 
the same quantities of any given chemical substances, undergoing 
the same chemical change under the same conditions, produce 
or absorb, according as the action is exothermal or endothermal, 
amounts of heat which are equal. 

The rate at which a given chemical action is allowed to take place 
has no influence on the total amount of heat consumed or produced. 
It may not at first sight appear obvious that rusting evolves heat, 
but a delicate thermometer will show that a heap of rusting nails is 
somewhat higher in temperature than surrounding bodies. Poor 
conductors, like oily rags and ill-dried hay, show a tendency to 
spontaneous combustion owing to accumulation of the slowly 
developing heat of oxidation (p. 37). The warmth of our own 
bodies is due to the same cause. 



176 COLLEGE CHEMISTRY 

It should be noted that production or absorption of heat is not, 
in itself, an evidence of chemical action. Physical changes are all 
likewise accompanied by the same phenomena. Thus, the evapo- 
ration of water absorbs heat, and condensation of a vapor and the 
crystallization of a supercooled liquid liberate heat. 

The heat of solution (c/. pp. 125, 145) is the heat liberated (or 
absorbed) on dissolving one mole of the substance in a large 
amount of water. A part of the water always undergoes chemical 
change (p. 139). The solute also frequently combines with a part 
of the water, or is ionized {q.v.), and the change in volume of the 
mixture (p. 138), as a physical phenomenon, would alone entail a 
heat-change. Hence this heat effect is partly chemical and partly 
physical in origin. 

Exercises. — 1. Which form of energy is delivered as such, 
and paid for as such, in most cities? 

2. How many calories are required to raise 500 g. of a substance 
of specific heat 0.5 from 15° to 37° (p. 174)? 

3. The combustion of 1 g. of sulphur to sulphur dioxide develops 
2220 calories. What is the heat of combustion of sulphur? Write 
the thermochemical equation. 



CHAPTER XIV 

CHEMICAL EQUILIBRIUM 

In spite of its formidable title, this chapter will introduce nothing 
novel. Its purpose is to collect together and organize more 
definitely a number of scattered facts and ideas which have already 
come up in various connections. On this account, however, it will 
be all the more necessary for the reader to refresh his remembrance 
of these facts and ideas by re-reading all pages to which reference 
is made. 

Reversible Actions. — In discussing Deacon's process (p. 156), 
it was stated that the action comes to rest although a large amount 
of both of the interacting substances (20 per cent at 345°) still 
remains available: .(20 per cent) 4HC1 + 2 <=* 2H 2 + 2C1 2 
(80 per cent). Now the materials thus left unused are presumably 
no less capable of interacting than were the parts which have 
already reacted. The solution of this mystery lies in the fact 
(p. 156) that the products themselves interact to reproduce the 
initial substances (read the equation backwards). Thus two 
changes, one of which undoes the work of the other, are going on 
simultaneously. In consequence of this, neither action can reach 
completion. As we should expect, experiment shows that it 
makes no difference whether we start with pure chlorine and steam, 
or with hydrogen chloride and oxygen; the proportions of the four 
substances found in the tube, after it has been kept at 345° for a 
sufficient time, are in both cases the same. A general statement 
may be founded on facts like this, to the effect that a chemical 
action must remain more or less incomplete when the reverse action 
also takes place under the same conditions. Two arrows pointing 
in opposite directions are used in equations representing reversible 
changes.* 

* The reader must avoid the idea that a reversible action is one which goes 
to completion, and then runs back to a certain extent. This conception would 
be contrary to the fact, and inexplicable by the kinetic method. 

177 



178 COLLEGE CHEMISTRY 

The foregoing example of a reversible action, and the following 
examples which very closely resemble it, should now be looked up 
and studied attentively. The discussion in this and the following 
sections, for which they furnish the basis, cannot otherwise be 
understood: (1) the interaction of chlorine and water (p. 161), 
which was fully discussed at the time; (2) the behavior of phos- 
phorus pentachloride vapor (p. 117); (3) the behavior of water 
vapor (p. 93), of phosphorus vapor (p. 117), of sulphur vapor 
(p. 117), and of iodine vapor (p. 117). 

When the action is one which is reversible, but, under the cir- 
cumstances being discussed, proceeds farther towards completion 
in one direction than in the other, the arrow will be modified to 
indicate this fact: 

Cl 2 + H 2 <=> HC1 + HCIO (p. 161). 

When this relative completeness is due to precipitation or vola- 
tilization, the fact may be indicated by vertical arrows: 

NaCl + H 2 S0 4 ^NaHS0 4 + HC1| (p. 141). 
NaCl j + H 2 S0 4 ±=>NaHS0 4 + HC1 (p. 143). 

Actions Which Proceed to Completion. — All chemical 
actions do not belong to the reversible, incomplete class. Many 
proceed uninterruptedly to exhaustion of one, or all, of the in- 
gredients. For example, equivalent amounts of magnesium and 
oxygen combine completely, 2Mg + 2 — > 2MgO. Here, how- 
ever, the product is not decomposed even at the white heat pro- 
duced by the vigor of the union. Indeed, magnesium oxide cannot 
be decomposed, and the action reversed, at any temperature we 
can command. The other complete actions, like the decomposi- 
tion of potassium chlorate (p. 27), are so because they are likewise 
irreversible. 

Explanation in Terms of Molecules. — Restating these facts 
in terms of the molecules will enable us to reason more clearly 
about this variety of chemical change. Suppose we start with the 
materials represented on one side only of such an equation, say the 
hydrogen chloride and oxygen in that on p. 177. The molecules 
of these materials will encounter one another frequently in the 
course of their movements. In- a certain proportion of these 



CHEMICAL EQUILIBRIUM 179 

collisions the chemical change will take place. In the earliest 
stages there will be few of the new kind of molecules (say, of 
chlorine and steam), but, as the action goes on, these will increase 
in number. There will be two consequences of this. In the first 
place the parent materials (in this case, hydrogen chloride and 
oxygen) will diminish in amount, the collisions between their 
molecules will become fewer, and the speed of the forward action 
will therefore become less and less. In the second place the in- 
crease in the number of molecules of the products will result in 
more frequent collisions betweenthem, in more frequent occurrence 
of the chemical change which they can undergo, and thus in an 
increase in the speed of the reverse action. The forward action 
begins at its maximum and decreases in speed progressively; the 
reverse action begins at zero and increases in speed. Finally the 
two speeds must become equal, and at that point perceptible change 
in the condition of the whole must cease (c/. pp. 88-89). 

The most immediate inference from this mode of viewing the 
matter is, that the apparent halt in the progress of the action 
does not indicate any cessation of either chemical change. Both 
changes must go on, in consequence of the continued encounters of 
the proper molecules. But since the two changes proceed with 
equal speeds they produce no alteration in the mass as a whole. In 
fact, the final state is one of equilibrium, and not of rest, one of 
balanced activity and not of repose. Hence, chemical changes 
which are reversible lead to that condition of seemingly suspended 
action which we speak of as chemical equilibrium. 

Chemical Equilibrium and its Characteristics. — The de- 
tailed discussion of the relations of liquid and vapor (pp. 78, 87- 
90), and of saturated solution and undissolved solid (pp. 127, 130- 
133), has already familiarized us with the term equilibrium and its 
significance. We can, in fact, apply to the discussion of any kind 
of reversible phenomena, the sets of ideas in regard to exchanges of 
molecules there elaborated. 

In particular, the reader will note that the three characteristics of 
a state of equilibrium, developed and illustrated in the case of the 
physical equilibrium between a liquid and its vapor (p. 89), apply 
also to a typical case of chemical equilibrium, such as that in 
Deacon's process now before us. Thus: 



180 COLLEGE CHEMISTRY 

1. There are the two opposing tendencies, which ultimately 
balance one another. Here they are the tendency of the steam and 
chlorine to produce hydrogen chloride and oxygen, and the tend- 
ency of the hydrogen chloride and oxygen to reproduce steam and 
chlorine by this interaction. In other words, they are the apparent 
activity of the hydrogen chloride and oxygen interaction, and the 
apparent activity * of the steam and chlorine reaction. 

2. At equilibrium the two opposing tendencies or activities are still 
in full operation, although their effects then neutralize one another. 

3 (and this is the chief mark of chemical, as it is of physical 
equilibrium). The system is in a sensitive state, so that a change 
in the conditions (temperature and pressure or concentration), even 
if slight, produces a corresponding change in the state of the system, 
and does this by favoring or disfavoring one of the two opposing 
tendencies or apparent activities. Such a change is called a dis- 
placement of the equilibrium, for the system settles down in a new 
state of equilibrium with new proportions of the two sets of sub- 
stances, corresponding to the changed conditions. Thus, in the 
present instance, a change from 345°, where there is 80 per cent 
of the material in the form of steam and chlorine, to 384° results 
in the diminution of this proportion to 75 per cent. The equilib- 
rium is affected by changes in concentration also, as we shall 
presently see (pp. 181, 186). 

Now, the foregoing facts show that the key to understanding 
chemical activities, their magnitudes, their changes, and especially 
their practical results, must lie in knowing how changes in the 
conditions affect them. Hence, to the chemist, familiarity with the 
influence of conditions on chemical phenomena must be of the 
greatest practical importance. We therefore address ourselves 
to the discussion of this subject. 

The "conditions" to be considered are familiar, — temperature, 
and concentration or, in the case of a gas, partial pressure. The 
"activity" of an action is accurately measured by the speed with 

* We use the term "apparent activity" for the activity as we see it. In 
the same action it varies with the conditions. The intrinsic activity or affinity, 
on the other hand, is the absolute activity of the action irrespective of condi- 
tions. Its value can be determined only by eliminating the effect of conditions, 
a matter which is too abstract for consideration here. The apparent activity 
is the practical thing which we observe. 



CHEMICAL EQUILIBRIUM 181 

which the action proceeds. Thus, if the foregoing section be re- 
examined, it will be seen that we spoke throughout of the speed, 
rather than of the tendency or activity. 

Finally, temperature and other conditions influence also the 
activities in, and therefore the speeds of, those actions which pro- 
ceed to completion, and are not reversible. Hence, unless our 
statements are expressly restricted to reversible actions and to 
states of equilibrium, they apply to all chemical changes. 

The Influence of Concentration. — In the first place, let us 
assume that the temperature is constant, and let us confine our 
attention for the present to the influence of concentration upon a 
chemical reaction. We have seen (p. 178) that the speed of a 
chemical change is determined by the frequency with which the 
molecules of the interacting substances encounter one another. 
The frequency of the encounters amongst a given set of molecules, 
resulting in a definite chemical change, will in turn evidently 
depend entirely upon the degree to which the molecules are con- 
centrated in each other's neighborhood. Larger amounts of one 
of the materials, for example, will not result in more rapid chemical 
action, if the larger amount of material is also scattered through a 
larger space. Chemical changes, therefore, are not accelerated by 
increasing the mere quantity of any ingredient, but only by in- 
creasing the concentration of its molecules. Thus, a large amount 
of hydrochloric acid with a piece of zinc will generate hydrogen no 
faster than a smaller amount. But substitution of more concen- 
trated acid will instantly increase the speed of the action. In the 
second case, the number of molecules of the acid reaching the zinc 
per second is greater, and this action, being non-reversible, pro- 
ceeds more rapidly to complete consumption of the zinc. So also, 
iron burns faster in oxygen (100 per cent) than in air (20 per cent 
oxygen). 

With a reversible action the effect on the speed is the same, ex- 
cepting that the continued activity of the reverse action prevents 
the direct one from reaching completion. 

Thus, if, in the action of hydrogen chloride upon oxygen, we 
introduce into the same space an extra amount of oxygen, this 
facilitates the formation of steam and chlorine by increasing the 
possibilities of encounter between molecules of hydrogen chloride 



182 COLLEGE CHEMISTRY 

and oxygen. At the same time it does not affect (c/. p. 72) the 
number of encounters in a given time of steam and chlorine mole- 
cules with one another which result in the reverse transformation. 
The proportion of chlorine (and steam) formed, therefore, from a 
given amount of hydrogen chloride will be greater, although the 
total possible (by complete consumption of the materials) has not 
been altered, since the quantity of one ingredient only has been 
increased. The introduction of an excess of hydrogen chloride 
would have had precisely the same effect. 

An Experimental Illustration. — A reaction in which the 
effects of different concentrations were carefully studied by Glad- 
stone (1855) affords a good illustration. If ferric chloride and 
ammonium thiocyanate are mixed in aqueous solution, a liquid 
containing the soluble, blood-red ferric thiocyanate is produced. 
The compound radicals are (NH 4 ) and (CNS), and the action is 
a simple double decomposition: 

FeCl 3 + 3 NH 4 CNS <=> Fe(CNS) 3 + 3 NH4CL 

The action is a reversible one, and the mixture is homogeneous, i.e., 
there is no precipitation. Now, if the two just-named salts are 
mixed in very dilute solution in the proportions required by the 
equation, say by adding 20 c.c. of a decinormal solution (p. 124) 
of each salt to several liters of water, a pale-reddish [solution is 
obtained. When this is divided into four parts, and one is kept for 
reference, the addition of a little of a concentrated solution of ferric 
chloride to one jar, and of ammonium thiocyanate to another, will 
be found to deepen the color by producing more of the ferric thio- 
cyanate. On the other hand, mixing a few drops of concentrated 
ammonium chloride solution with the fourth portion will be found 
to remove the color almost entirely, on account of its influence in 
accelerating the backward change. 

The Law of Molecular Concentration. — The general prin- 
ciple discussed and illustrated in this section may be called the law 
of molecular concentration, and may be stated as follows: In 
every chemical change the apparent activity, and therefore the speed 
of the action, is proportional to the molecular concentration of each 
interacting substance. This holds whether the action is reversible 
or not. 



CHEMICAL EQUILIBRIUM 183 

We shall next give a more precise, semi-mathematical formula- 
tion of this law, as this formulation will be of use later,* and then 
proceed to illustrate the application of the law, by showing how 
it explains large classes of actions of which we have already en- 
countered many examples. 

^Formulation of the Law of Molecular Concentration. — 

The mathematical formulation of the law describing the influence 
of the concentration of the molecules of each participating sub- 
stance upon the speed of the action is extremely simple. When the 
actual concentrations of the molecules are specified (in moles, 
pp. 102, 125, per liter), and the speed is suitably expressed (in moles 
transformed per minute or per hour), we find that the speed is 
proportional to the concentration of each molecule appearing in 
the molecular equation for the action. Thus in the interaction of 
hydrochloric and hypochlorous acids (the reverse of the action of 
chlorine on water, pp. 161, 178), if [HC1] and [HCIO] represent the 
concentrations of the molecules HC1 and HCIO, and A; is a con- 
stant, and S is the speed, then 

[HC1] X [HCIO] X k = S. 

Again, for the dissociation of phosphorus pentachloride vapor 
into phosphorus trichloride and chlorine (p. 117) : PCI5 — > PC1 3 + 
Cl 2 , if [PCI5] represents the concentration of the PC1 5 molecules, ki 
is a constant, and Si is the speed of decomposition: 

[PC1J X h = Si. 

Similarly, for the reverse action : PC1 3 + Cl 2 — > PC1 5 , if [PC1 3 ] and 
[Cl 2 ] stand for the molecular concentrations of these substances: 

[PCIJ X [C1J Xk 2 = S 2 . 

The constant has a different value in each separate action. It 
includes the value of the intrinsic affinity or activity of the sub- 
stances, and the catalytic effect (p. 29), if any, of the materials 
present. 

* This formulation of the law is not required, or referred to in the sections 
which follow. The section and the following one may therefore be omitted 
for the present and be taken up in connection with Chaps. XX and XXXV. 



184 COLLEGE CHEMISTRY 

Formulation of the Condition for Chemical Equilibrium. 

— The foregoing plan may be used further to formulate the con- 
dition for chemical equilibrium. As we have seen (p. 179), a 
characteristic of a system in chemical equilibrium is that the speeds 
of the forward and reverse actions have become equal. If, then, 
[PCl 5 ]eqm., [PCVIeqm., and [Cl 2 ] e q m . now represent the molecular con- 
centrations when the system has reached equilibrium, then, since 
the speeds are equal: 

[PClJeqm. X [Cl 2 ]eqm. X fe = [PClJeqm. X fe, 
[PCl 3 WX[Cl 2 ]eq m , h = congtant . 
L-rUl5jeqm. #2 

In words, this means that if we change the amount of the penta- 
chloride placed in the vessel, or if we use amounts of chlorine and 
trichloride which are not equivalent, the numerical value at equilib- 
rium of each concentration ([PC1 3 ] etc.) will, of course, be differ- 
ent, but the product of the concentrations of trichloride and 
chlorine, divided by the concentration of the pentachloride, will 
always give the same numerical value for the constant at the same 
temperature. This numerical value is called the equilibrium 
constant. 

Applications: The Forward Action. Homogeneous and 
Inhomogeneous Systems. — While there are all degrees of speed 
in chemical actions, yet in practice we quickly distinguish two 
different classes. There is a class of actions of which most exam- 
ples are almost instantaneously accomplished, and a class in which, 
frequently, the operation takes minutes or even hours. The 
classes overlap, but, in a general way, the following distinction may 
be made. 

To the former, speedy class belong the explosion of hydrogen 
and oxygen or other gaseous mixtures, and the interactions when 
solutions are mixed, as in precipitations. In view of the foregoing 
explanations, we perceive that the rapid accomplishment of such 
actions is due, not so much to any especially great intrinsic affinity, 
as to the homogeneous state of mixture of the interacting materials. 
This, of course, is a purely physical, and not a chemical motive for 
speedy interaction. In intimate mixtures, every molecule has an 
equal opportunity freely to encounter every other molecule and 



CHEMICAL EQUILIBRIUM 185 

there is therefore no mechanical impediment to the operation of 
the affinities of the substances. Hence the apparent activity is 
great. 

To the second class, comprising the slower actions, belong cases 
like the interaction of a piece of zinc with hydrochloric .acid, or of 
manganese dioxide (p. 158) with the same acid, whereby hydrogen 
and chlorine, respectively, are slowly evolved, and the solid is grad- 
ually consumed. Here the hindrance is evidently the fact that 
the interacting substances are not intimately mixed. In the slow 
actions, the system is inhomogeneous. Pulverizing the solid before 
use will increase the speed, indeed, by providing more surface and 
better mutual contact, but will not transfer the action to the rapid 
class. It is chiefly the dissolved part of the substance which inter- 
acts, for chemical action takes place between molecules, and only 
the dissolved part is disintegrated in such a way that the molecules 
are readily accessible. Thus, the action is held back by continual 
waiting for the slow replenishment, from the " insoluble" solid, of 
the supply of dissolved molecules. In the cases cited, the restrain- 
ing influence of the dissolving process, which is part of the whole 
phenomenon, may be formulated thus: 

Zn(solid) 1=5 Zn(dslvd) + 2HC1 -> ZnCl 2 + H 2 . 
Mn0 2 (solid) <=? Mn0 2 (dslvd) + 4HC1 -> MnCl 2 + 2H 2 + Cl 2 . 

Here, again, the mechanical details, depending on physical prop- 
erties, have more to do with the progress of the action than has 
the chemical affinity. In terms of the law of concentration, the 
action is slow, and the apparent activity small, because the con- 
centration of the acting molecules of one of the substances is very 
small, and cannot be increased because of low solubility. 

Applications: The Reverse Action. Displacement of 
Equilibria. — We have seen (p. 182) that one way in which a re- 
versible action may be forced nearer to completion, in one direction 
or the other, is the introduction of an excess of one of the ingre- 
dients contributing to the forward action. This method of dis- 
placing the equilibrium point, however, cannot be very effective, 
unless it is possible to introduce an exceedingly large excess of the 
selected ingredient in a high degree of molecular concentration, 
since this operation does not in any way affect or, in particular, 




186 COLLEGE CHEMISTRY 

restrain the reverse action which is continually undoing the work of 
the forward one. A much more effective means of furthering the 
desired direction of such actions is found, therefore, in the restraint 
or practical annulment of the reverse action. A good way of 
accomplishing this is to allow the products of the direct action to 
separate into an inhomogeneous mixture. Any agency which could 
remove the water vapor as fast as it was formed by the interac- 
tion of hydrogen chloride and oxygen, for example, would entirely 
stop the reproduction of these substances, and so would enable 
the forward action (4HC1 + 2 — > 2H 2 + Cl 2 ) to run to completion. 
This might be realized by causing one end of a sealed tube 
charged with the substances, after the contents had settled down 

to a condition of equilibrium, to 
project from the bath in which the 
whole had been kept at 345° (Fig. 
74, which is simply diagrammatic). 
By cooling this end, a large part 
of the steam would quickly be con- 
densed in it to the liquid form, 
while the other substances would remain gaseous. In other words, 
the concentration of the water vapor would be greatly reduced. In 
fact, only the trace of vapor which cold water gives would then be 
available to interact with the chlorine, and reproduce hydrogen chlo- 
ride. Meanwhile the decomposition of the latter would go on, and 
thus, eventually, almost all the water would be found in one end of 
the tube, and the chlorine, all free, would occupy the rest. By this 
purely mechanical adjustment the chemical change would therefore 
be carried from 80 per cent completion to almost absolute completion: 
4HC1 + 2 *± 2C1 2 + 2H 2 (vapor) t± 2H 2 (liq.) 
If, on the other hand, arrangements were made to have pow- 
dered marble, in a sealed bulb of thin glass, enclosed in the tube, 
we might imagine the very opposite of the above effect to be pro- 
duced. The breaking of the bulb of marble, when equilibrium 
had been reached, would provide means for the removal of all the 
hydrogen chloride,* while the other three substances would still be 

* The hydrogen chloride would be destroyed by interaction with the marble: 
2HC1 + CaCOs -> CaCl 2 + C0 2 + H 2 0. 
The calcium chloride is a solid. The gas, carbon dioxide, does not interact 
with the other substances, and would not, therefore, interfere with the forma- 
tion of fresh hydrogen chloride. 



CHEMICAL EQUILIBRIUM 187 

gaseous. Thus, the compound (HC1) having been reduced in 
concentration to the point of being removed entirely, there would 
be no direct action to undo the work of the reverse action. The 
whole chlorine would, therefore, soon have passed through the 
form HC1. Hence, by another mechanical arrangement, an action 
which ordinarily could progress to only 20 per cent would be 
turned into a complete one: 

2C1 2 + 2H 2 <=» 2H 2 + 4HC1 (+CaC0 3 -» CaCl 2 + H 2 + C0 2 ). 

Reversibility Usually Avoided, — In every-day chemical 
work, since our object is usually to prepare some one substance, 
chemists either avoid chemical changes which are notably re- 
versible, or adjust the conditions, as is done in the foregoing 
illustrations, so that the reverse of the action which they desire is 
prevented. In consequence of this, when carrying out the direc- 
tions for making familiar preparations, the fact that such actions 
are reversible at all very readily escapes our notice. Arranging 
the conditions so that the separation of a solid body by precipita- 
tion, or the liberation of a gas, takes place, are the two commonest 
ways of rendering a reversible action complete. Excellent ex- 
amples of both of these are furnished by the chemical change 
used in producing hydrogen chloride by the interaction of salt and 
sulphuric acid, the full discussion of which (p. 142) should now be 
studied attentively in the light of these explanations. 

History. — The conceptions discussed in this chapter are not 
new, although they have come into general use rather recently. 
The law of reaction speed, and the influence of the concentrations 
of the reacting substance thereon (p. 183), was set forth and 
formulated by Wilhelmy as early as 1850. Gladstone (1855) 
studied quantitatively the influence of concentration in cases of 
chemical equilibrium (p. 182). The kinetic explanation (p. 178) 
was developed by Williamson (1851) . Finally, the laws of chemical 
equilibrium were formulated more explicitly and applied more 
thoroughly by two Swedish chemists, Guldberg and Waage 
(1864-9). 

The Influence of Temperature on the Speed of any Re- 
action. — The activity of chemical change, and therefore the 



188 COLLEGE CHEMISTRY 

speed of all chemical changes, is increased by raising the temperature 

and diminished by lowering it (c/. p. 59). Thus, zinc displaces 
hydrogen more rapidly from hot than from cold hydrochloric acid. 
Different actions are affected in different degrees, and no simple 
rule accurately defining the effect can be given. Roughly speak- 
ing, however, a rise of 10° doubles the speed of every action. A 
rise of 100° will therefore make the speed roughly 1024 times 
greater. Hence, when the chemist finds that two substances 
show no evidence of interaction, he infers that there must be either 
slow action or none, and he seeks to settle the question quickly by 
heating the mixture. 

The Influence of Temperature on a System in Equilib- 
rium. — In a reversible change the two opposing reactions are 
different actions and their speeds are therefore affected in different 
degrees by the same alteration in temperature. Hence, when the 
temperature is changed, the relative amount of the two sets of 
materials present is~altered and the equilibrium is displaced. Thus, 
in Deacon's process, a rise of 40° in the temperature displaces the 
equilibrium backwards (p. 180), and diminishes the yield of chlo- 
rine by 5 per cent. In the vapor of phosphorus pentachloride 
(p. 117), the displacement is in the opposite direction. The vapor 
is a mixture of the pentachloride with the trichloride and free 
chlorine: PC1 5 ^ PC1 3 + Cl 2 . At 200°, 51.5 per cent of the 
material is present as pentachloride and 48.5 per cent as trichloride 
and chlorine. Raising the temperature to 250° changes the pro- 
portions to 20 per cent and 80 per cent, respectively. At 300° only 
3 per cent of the pentachloride remains. Evidently, here, raising 
the temperature favors the decomposition of the pentachloride, 
and therefore increases the speed of its dissociation more than it 
does the speed of the reunion of the trichloride and chlorine. 

Van't Hojf's Law. — One use of a law is to enable us to answer 
a question, when we have not in memory the fact constituting the 
answer, and even when we have never read or heard the fact. The 
law or rule enables a little reasoning to take the place of a vast 
amount of memorizing. Thus, to answer the question: Does 
sodium chloride always have the same composition, it is not 
necessary to have read and to remember all, or any of the numerous 



CHEMICAL EQUILIBRIUM 189 

investigations of this substance that have been made. We simply 
refer the question, mentally, to the law of definite proportions, 
and say "yes." Now the facts mentioned above are connected 
by a law which will answer many practical questions in chemistry. 
When phosphorus trichloride and chlorine combine (to form 
PC1 5 ), heat is given out. Conversely, when phosphorus penta- 
chloride dissociates, heat is absorbed: 

PC1 5 + 30,000 cal. <=> PCI3 + CI* 

Now, when the temperature is raised, the action proceeds in the 
direction of decomposing more of the pentachloride. That is, the 
equilibrium is displaced in the direction which absorbs heat. 

In Deacon's process, we find that the interaction of hydrogen 
chloride and oxygen liberates heat, 

4HC1 + 2 <=± 2H 2 + 2C1 2 + 28,000 cal., 

and in this action raising the temperature drives the equilibrium 
backwards, and a lowering in the temperature is required to increase 
the yield of chlorine. 

The rule is obvious, and applies to all reversible reactions: 
When the temperature of a system in equilibrium is raised, the equi- 
librium point is displaced in the direction which absorbs heat. In 
other words, a rise in temperature favors the interaction of that 
one of the two sets of materials to which the heat is added (+ sign) 
in the equation. If the equation happens to be written with a 
negative heat of reaction (e.g., p. 175), the heat can, of course, be 
transferred to the other side with its sign changed. This law is 
known as Van't Hoff's law of mobile equilibrium. 

This law is of practical value. More than once, in chemical 
factories, much time and money have been spent on trying to 
arrange machinery to give a better yield of some substance at a 
high temperature, when a reference to this law would have shown 
that the chief change necessary was to use a lower temperature. 
We shall frequently have occasion to refer to this law. 

Application to Physical Equilibria. — Van't Hoff's law 
applies also to physical processes. Thus, as the temperature rises, 
a substance which absorbs heat in dissolving will become more 
soluble. This is the commoner case, as is shown by the way in 



190 COLLEGE CHEMISTRY 

which most solubility curves (Fig. 58, p. 131) ascend with rising 
temperature. Conversely, a substance which gives out heat in 
dissolving is less soluble with rising temperature in a solution 
already almost saturated with the compound. For example, anhy- 
drous sodium sulphate gives out heat in dissolving, and so its 
solubility diminishes (Fig. 59, p. 132), with rising temperature. 

Again, the vaporization of a liquid absorbs heat, and so an in- 
crease in temperature will increase the pressure, and therefore the 
concentration of its vapor (p. 87). 

Le Chatelier's Law. — The above mentioned law is really a 
particular case of a more general one. If some stress (e.g., by 
change of temperature, pressure, or concentration) is brought to 
bear on a system in equilibrium, the equilibrium is displaced in the 
direction which tends to undo the effect of the stress. Thus, raising 
the temperature furthers the change which absorbs heat — and 
therefore would tend to lower the temperature. Increasing the 
concentration of the molecules pushes the action in the direction 
which uses up these very molecules (p. 181). Again pressure 
causes ice to melt, because the water which is formed occupies a 
smaller volume, and this change tends to relieve the pressure. 
But pressure will not cause most substances to melt, because 
usually the liquid form occupies a greater volume and its produc- 
tion would tend to increase pressure. 

Summary. — In this chapter we have answered three ques- 
tions: 

1. Why do some chemical actions cease, while still incomplete? 
Answer: They are reversible. 

2. What explains the position of the equilibrium point? An- 
swers: (a) Equal effects of opposed molecular actions; (6) Equality 
in speed of opposed reactions. 

3. What will displace the equilibrium point? Answer: (a) 
Change in concentration of one (or more) of the substances; (6) 
Change in the temperature. 

Exercises. — 1. Explain the completeness of the action by 
which hydrogen chloride and water, respectively, are formed by 
direct union of the elements. 



CHEMICAL EQUILIBRIUM 191 

2. Explain the completeness of the action by which silver 
chloride (p. 148) is formed. 

3. Explain why the decomposition of potassium chlorate is 
complete. 

4. In view of the statement on p. 14, explain why mercuric oxide 
is completely decomposed by heating. Point out the resemblance 
between this experiment and the imaginary one illustrated in 
Fig. 74 (p. 186). 

5. Why can magnetic oxide of iron be reduced completely by 
a stream of hydrogen (p. 59), and iron oxidized completely by a 
current of steam (p. 51)? 

6. With the phosphorus pentachloride system, say at 250°, what 
effect would suddenly enlarging the space containing a given 
amount of the vapor produce? What would be the effect of di- 
minishing the space? What would be the effect of introducing 
additional chlorine into the same space (p. 181)? 

7. By what practical means could the degree of dissociation of 
sulphur vapor (S 8 ) be reduced, without changing the temperature 
(P- H7)? 

8. What inference should you draw from the fact that: (a) the 
solubilities of potassium nitrate and of Glauber's salt (p. 132) 
increase with rise in temperature; (6) that those of calcium hy- 
droxide (p. 130) and triethylamine decrease with rise in tempera- 
ture? 



CHAPTER XV 

THE HALOGEN FAMILY 

The elements to which we have so far devoted most attention 
have been oxygen, hydrogen, and chlorine. If we recall the chemi- 
cal properties and relations of these elements we shall recognize 
the fact that they all possess very distinct individualities. 

The Chemical Relations of Elements. — Hydrogen is a sub- 
stance (p. 58) which unites readily with oxygen and chlorine, 
less readily with other non-metals, and scarcely at all with metals. 
Oxygen and chlorine resemble each other somewhat in the great- 
ness of their chemical activity and the variety of free elements with 
which they are capable of uniting, but differ markedly in what we 
have called their chemical relations (p. 163). The resulting com- 
pounds belong, in fact, to quite different classes — oxygen forms 
oxides, chlorine forms chlorides — and elements are considered 
similar only when they resemble each other in chemical relations, 
and produce, by combination with the same element, compounds 
having similar chemical properties. Thus, the common oxide of 
hydrogen, water, is a neutral substance, and is chemically rather in- 
different. The chloride of hydrogen in aqueous solution is a strong 
acid and is chemically very active.* If all the other chemical ele- 
ments differed from one another as much as do these three, they 
would be incapable of classification. In reality, however, we find 
that the elements can be grouped together in sets. They are classi- 
fied according to the kind of substances with which they combine 
and the chemical nature of the products. In some families the re- 
semblance is close, in others less close. The present group is of 
the former class, and will serve, therefore, as a convenient begin- 

* The difference between oxides and chlorides is seen in their behavior. 
Thus, oxides often unite with water to form acids or bases (p. 94). Chlorides 
do not unite with water to form new substances with marked characteristics 
(cf. p. 96). 

192 



BROMINE 193 

ning in the work of tracing relations between the elements and in 
classifying the facts of descriptive chemistry. 

The Chemical Relations of the Halogens. — The bromide 
(NaBr), iodide (Nal), and, to a less extent, the fluoride (NaF) of 
sodium, resemble sodium chloride (NaCl) in appearance and be- 
havior. From this fact, chlorine, bromine, iodine, and fluorine are 
known as the halogens (Gk., salt producers), and their compounds 
are named the halides. The halogens, as the above formulae show, 
are univalent. They all form compounds with hydrogen, and 
these compounds closely resemble hydrogen chloride. For ex- 
ample, they are colorless, they are gases (except hydrogen fluoride, 
a very volatile liquid), they are very soluble in water, and their 
solutions are acids. Other relations will be given in a summary 
at the end of the chapter. 

Bromine Br 2 

Occurrence. — The compounds of chlorine, bromine, and 
iodine usually occur together in nature, while the compounds of 
fluorine are not found in the same sources. Bromine occurs 
chiefly in the form of the bromides of sodium and magnesium, in 
the upper laj^ers of the natural beds of rock salt. Liebig made it 
from this source and a little later Bal- 
lard (1826) made it also and recognized 
it as a new element. 




Preparation. — In the chemical 
point of view there are three distinct Fia 75 - 

ways in which bromine is made. 1. The first of these is closely 
related to the common method of preparing chlorine (p. 158). 
As hydrobromic acid, unlike hydrochloric acid, is not formed ex- 
tensively in connection with any chemical industry, potassium 
bromide is treated in a retort (Fig. 75) with concentrated sul- 
phuric acid, and the product is oxidized with powdered manganese 
dioxide in one operation. (For equation see next section.) 
Bromine being a volatile liquid, while the sulphates of potassium 
and manganese are involatile, its vapor passes off when the above 
mixture is heated. It is condensed in a flask surrounded by cold 
water. 



194 COLLEGE CHEMISTRY 

2. The second method of preparing bromine depends on the fact 
that chlorine is a more active element and displaces bromine from 
combination. When, therefore, chlorine is passed into a solution 
of potassium or sodium bromide, potassium or sodium chloride is 
formed and the bromine liberated: 

2NaBr + Cl 2 -> 2NaCl + Br 2 . 

When the liquid is warmed, the bromine passes off along with a 
part of the water, and may be condensed as before. 

3. Aqueous solutions of soluble bromides may be decomposed 
by means of a current of electricity. The bromine is set free at the 
positive electrode. 

Commercial Extraction. — Two-thirds of the world's supply 
is obtained from Stassfurt, where, after the extraction of the 
potassium chloride from the impure carnallite (KCl,MgCl 2 ,6H 2 0), 
the mother-liquor is found to contain the more soluble sodium 
and magnesium bromides in considerable quantities. The warm 
mother-liquor trickles down over round stones in a tower. The 
chlorine is introduced from below and dissolves in the liquid. The 
bromine is thus liberated and passes off as vapor. A part of our 
supply of bromine is obtained from the brines of Ohio, West 
Virginia, and Kentucky, from which, after most of the common 
salt has been removed by crystallization, the bromine is obtained 
by the first method. In Michigan the brines are treated with 
electrolytic chlorine by the second method. 

Partial Equations, a Plan for Making Complex Equations. 

— When an equation involves more than two initial substances or 
products, as does the one for the first method of preparing bromine, 
it cannot readily be worked out by the method formerly recom- 
mended (p. 51). After the formulae of all the substances, on both 
sides, have been set down, it is difficult to hit upon the proper co- 
efficients required to balance the equation. In such cases, a good 
plan is to select two of the initial substances, and make a partial 
equation showing part of the action and including at least one 
actual product. Any unused units (not constituting a product) 
are then set down also and treated as a balance. Thus the first 



BROMINE 195 

two of the substances named will furnish potassium-hydrogen 
sulphate : 

Partial, 1 : KBr + H 2 S0 4 -» KHS0 4 (+ HBr) . (1) 

Similarly, the manganese dioxide and sulphuric acid will give 
manganous sulphate: 

Partial, 2: MnO a + H 2 S0 4 -> MnS0 4 + H 2 (+ 0). (2) 

We then perceive that the bromine must come from the oxidation 
of the first balance (HBr) by the second (O) : 

Partial, 3: (2HBr) + (0) -> H 2 + Br 2 . (3) 

The third partial equation shows that 2HBr will be needed for the 
amount of obtainable from Mn0 2 , so we go back to (1) and 
multiply it by two throughout : 

2KBr + 2H 2 S0 4 -> 2KHS0 4 (+ 2HBr). (1) 

MnOa + H 2 S0 4 -> MnS0 4 + H 2 (+ 0). (2) 

(2HBr) + (0) -» H 2 + Br 2 . (3) 

2KBr + 3H 2 S0 4 -\ Mn0 2 -+ 2KHS0 4 + MnS0 4 + 2H 2 + Br 2 . 

When we now add the real substances used and produced, as they 
occur in these partial equations, and leave out the balances, which 
have been adjusted so as to cancel one another, we obtain the final 
equation for the action. It must be observed that this subdivi- 
sion of the action into parts is a purely arithmetical device. It 
is still true, however, that we are aided in the selection of partial 
actions at each step by following some plausible theory as to 
stages for the action which would be chemically conceivable. 

Physical Properties. — Bromine is a dark-red liquid (sp. gr. 
3.18). It boils at 59°, forming a deep-red vapor, and even at ordi- 
nary temperatures gives a high vapor pressure (150 mm. at 18°) and 
evaporates quickly. When cooled it forms red, needle-shaped 
crystals (m.-p. — 7.3°). A saturated aqueous solution (bromine- 
water) contains 3 parts of bromine in 100 parts of water. The 
element is much more soluble in carbon disulphide, alcohol, and 
other organic solvents. Up to 750°, the G.M.V. weighs 160 g. 
(corresponding to Br 2 ), against 28.955 g. for air. 

Bromine (Gk., a stench) has a most pungent odor. It has a 



196 COLLEGE CHEMISTRY 

very irritating effect on the mucous membrane of the nostrils and 
throat. If spilled upon the hands it destroys the tissues and leaves 
sores which are liable to infection. 

Free bromine has no effect upon starch emulsion (see Iodine). 

Chemical Properties. — A jet of hydrogen gas burns in 
bromine vapor. The union is much slower than in the case of 
chlorine (Heat of formation, + 12,300 cal.). 

Bromine forms compounds directly, both with non-metals, like 
phosphorus and arsenic, and with most of the metals, which catch 
fire when thrown into the vapor. In all cases the interaction is 
less violent than when chlorine is used, and bromine is displaced 
from combination with hydrogen and with the metals by free 
chlorine. 

Silver bromide is the sensitive material in photographic plates, 
and potassium and sodium bromides are used as sedatives in 
medicine. Bromine is employed in the preparation of organic 
dyes. 

Hydrogen Bromide HBr 

Preparation. — It might be expected that the most convenient 
way of producing this compound would be similar to that used in 
preparing hydrogen chloride, namely, by the action of concentrated 
sulphuric acid upon some common bromide, such as potassium 
bromide (KBr + H 2 S0 4 <=± HBr + KHS0 4 ). Hydrogen bromide 
being less stable, however, a large part of it is oxidized by the 
sulphuric acid and the product is mixed with sulphur dioxide and 
free bromine. 

H 2 S0 4 + 2HBr -> 2H 2 + S0 2 T + Br 2 1 . 

Since all acids decompose all salts more or less, use of an acid 
which does not give up its oxygen so readily, such as phosphoric 
acid, will yield pure hydrogen bromide (KBr + H3PO4 — > HBr f 
+ KH 2 P0 4 ). The small solubility of the salt in concentrated 
phosphoric acid retards the interaction and makes the evolution of 
the gas very slow, however. 

Pure hydrogen bromide is best prepared by the action of water 
upon phosphorus tribromide (see Hydrolysis, below). When 
bromine and phosphorus are mixed, a violent union of the two 



HYDROGEN BROMIDE 



197 





Br 


H 


/ 


Br + H 


\ 


Br 


H 



elements takes place, producing phosphorus tribromide PBr 3 . 
This substance, which is a colorless liquid, is in turn broken up 
with great ease by water, producing phosphorous acid, which is not 
volatile, and gaseous hydrogen bromide: 

OH /OH 

OH->3HBr + P-OH 
OH x OH 

In practice, these two actions are carried on simultaneously. To 
diminish the vigor of the interaction, red phosphorus is taken in- 
stead of yellow, and is mixed with two or three times its weight of 
sand in a flask (Fig. 76) . A small quantity of water is added. Ex- 
cess of water must be avoided, as the hydrogen bromide produced 
is extremely soluble, and would there- 
fore be retained in the flask instead of 
being disengaged as gas. The bromine 
is placed in the dropping funnel, and 
admitted, a little at a time, to the 
mixture. The gas produced is passed 
through a U-tube containing red phos- 
phorus mixed with glass beads. The 
phosphorus combines with any free 
bromine carried along with the gas. 
The second U-tube, containing water, may be attached when a 
solution of the gas is required. The gas may be collected in a 
jar by upward displacement of air. 

Hydrolysis, — The interaction of water with phosphorus tri- 
bromide (foregoing section) illustrates an important property of 
water (p. 92) . The action is a double decomposition in which water 
is one of the interacting substances and is called an hydrolysis (Gk., 
loosening by water). The water divides into the radicals H and 
OH, and the former unites with the more active non-metal in the 
substance (the bromine, in PBr 3 ) and the hydroxyl with the other 
element. For example, PC1 3 + 3HOH -> P(OH) 3 + 3HC1. All 
the halides of the non-metals are thus hydrolyzed, as are also some 
other classes of compounds. 

Physical Properties. — Hydrogen bromide is a colorless gas 
with a sharp odor. It is two and a half times as heavy as air. It is 




Fig. 76. 



198 COLLEGE CHEMISTRY 

easily reduced to the liquid condition (b.-p. — 69°). It is ex- 
ceedingly soluble in water, and in contact with moist air condenses 
the water vapor to clouds of liquid particles. Pure hydrogen 
bromide, whether in the gaseous condition or in the liquefied form, 
is a nonconductor of electricity (see below). 

Chemical Properties, — The properties are like those of 
hydrogen chloride (p. 145). It is somewhat less stable, and dis- 
sociation begins to be noticeable at 800°. When free from water, 
it is not an acid (see below). The gas interacts vigorously with 
chlorine, hydrogen chloride and free bromine being produced, 
2HBr + Cl 2 -> 2HC1 + Br 2 . What are the relative volumes (p. 
150)? 

Chemical Properties of Hydrobromic Acid HBr, Aq. — The 

solution of the hydrogen bromide in water is an active acid (cf. 
p. 52). It conducts electricity extremely well. In contact with 
certain metals, and with oxides of metals and hydroxides of metals, 
it behaves exactly like hydrochloric acid (p. 146) . In the first case, 
hydrogen is set free and the bromide of the metal produced. In 
the other two cases, water and the bromides of the metals are 
produced. For example: Zn(OH) 2 + 2HBr -» ZnBr 2 + 2H 2 0. 
Oxidizing agents set bromine free from hydrobromic acid, even 
sulphuric acid, which does not act upon hydrochloric acid, being 
able to do this (p. 196). Chlorine dissolved in water displaces 
bromine from hydrobromic acid and from soluble bromides with 
ease (test for bromides) . 

Iodine I 2 

Occurrence, — Iodine occurs in sea-water, about one-fifth of 
it in algae and four-fifths in organic compounds. Certain species 
of sea-weed, known in Scotland as kelp and in Normandy as varec, 
remove it from the water. The ash of the sea-weed sometimes 
contains as much as two per cent, or even more. The other chief 
source of iodine is in Chile saltpeter (mainly NaNOs), in which it 
is present in the form of about 0.2 per cent of sodium iodate 
NaI0 3 and sodium iodide. Most of the iodine of commerce is 
obtained from this source and only a little from sea-weed. The 
largest proportion of iodine in the human body is in the thyroid 



IODINE 199 

gland. In diseases like goitre and cretinism, where the thyroid 
is ill-developed, injection of a substance called iodothyrine, ex- 
tracted from sheep's thyroids, produces marked improvement. 

Preparation. — 1. In factories where the iodine is extracted 
from sea- weed, the latter is carbonized in retorts and sodium iodide 
is extracted with water from the residue. This is then treated 
with manganese dioxide and sulphuric acid. The quantity of 
manganese dioxide is carefully measured so as to be just sufficient 
to set free the iodine contained in the liquid, without proceeding 
farther to the liberation of the chlorine which it contains in much 
larger amounts. When the mixture is heated, the iodine passes off 
in the form of vapor, and is condensed in a suitable receiver. 
The action (c/. pp. 157, 194) is: 

2NaI + Mn0 2 + 3H 2 S0 4 -» MnS0 4 + 2NaHS0 4 + 2H 2 + I 2 . 

2. In France the treatment is similar, excepting that chlorine is 
used to liberate the iodine in the last stage (2NaI+Cl 2 — >2NaCl+I 2 ). 
The quantity is adjusted so that excess may not be employed. 
The iodine, being insoluble, forms a dense precipitate and, when 
the liquid is pressed out, it remains behind in the form of a 
paste. 

3. Electricity could also be used for the decomposition of this 
mother-liquor. The iodine is set free at the positive electrode. 

In all cases the iodine is purified by distillation with a little 
powdered potassium iodide. It condenses in the solid form di- 
rectly, in glittering, black plates (sublimed iodine). The distilla- 
tion of a solid body, when a condensation takes place directly to 
the solid form, is spoken of as sublimation. 

Physical Properties. — Iodine (Gk., like a violet) is a black, 
solid substance (sp. gr. 5), exhibiting large crystalline plates of 
rhombic form. It melts at 114°, and boils at 184°. The vapor has 
at first a reddish-violet tint, and on being more strongly heated 
becomes deep blue (see next section). 

Iodine is very slightly soluble in water (about 1 : 6000), and the 
solution has a scarcely perceptible brown tint. It is much more 
soluble in carbon disulphide (p. 12) and in chloroform, in which 
it gives violet solutions. In alcohol it gives a solution which is 



200 COLLEGE CHEMISTRY 

brown, the iodine being in a condition of feeble combination, and 
not simply in solution. An aqueous solution of potassium iodide, 
hydrogen iodide, or any other iodide, has likewise the power to 
take up large quantities of iodine. Here the formation of definite 
compounds (such as, KI + I 2 ^±KI 3 ), by a reversible action, 
accounts for the amount of iodine taken up. 

The behavior of free iodine towards starch forms a distinctive 
test for both substances (c/. p. 3). The pale-brown aqueous 
solution, for example, when added to starch emulsion, produces a 
deep-blue color. This blue substance is not a chemical compound. 
The iodine is adsorbed by the starch, which is in colloidal suspension 
(q.v.). 

Chemical Properties. — The molecular weight of iodine, ascer- 
tained by weighing the vapor at temperatures from the boiling- 
point up to 700°, is 253.8. The atomic weight being 126.92, the 
molecule contains two atoms. Beyond 700°, the vapor diminishes 
in density more rapidly than Charles' law would lead us to expect, 
and at 1700° the molecular weight has fallen to 127 (cf. p. 117). 
As the vapor is heated, a larger and larger proportion of the mole- 
cules is broken up, until the decomposition has become complete. 
As in all cases of dissociation, when the vapor is cooled the atoms 
recombine to form molecules. This is the most notable case in 
which we encounter both the monatomic and the diatomic forms 
of the same element. The heat given out when the atoms reunite 
to form the molecules is very considerable (21 ?=± I 2 + 28,500 cal.), 
indicating that the chemical union of two atoms of identical nature 
may be as vigorous as that of two atoms of different chemical 
substances. The heat of union of atomic hydrogen (p. 113) is 
even greater (2H <=± H 2 + 90,000 cal.). In both cases, in accord- 
ance with Van't Hoff's law (p. 189), raising the temperature 
increases the dissociation, because that is the direction in which 
heat is absorbed. 

Iodine unites very slowly with hydrogen, even when heated. 
It unites directly with some non-metals and with the majority of 
the metals. When phosphorus is presented in the yellow form, 
the action takes place spontaneously without the assistance of 
heat. Both chlorine and bromine displace iodine from combina- 
tion with hydrogen and the metals (2HI + Br 2 — * 2HBr + I 2 ). 



HYDROGEN IODIDE 201 

The action may be brought about either with the substances in 
dry form or with their aqueous solutions. 

Iodine and its compounds are much used in the arts and medicine. 
Iodine is applied, in the form of an alcoholic solution (tincture of 
iodine), for the reduction of some swellings. It is required in 
making iodoform CHI 3 , and the iodides of potassium, rubidium, 
and sodium, which are used in medicine. The emulsion used in 
making photographic dry-plates contains silver iodide Agl. 

Hydrogen Iodide HI 

Preparation, — The direct union of hydrogen and iodine can- 
not be employed in preparing pure hydrogen iodide (see below). 

The action of concentrated sulphuric acid upon potassium iodide 
is equally inapplicable. In this case, as in that of hydrogen 
bromide (p. 196), the sulphuric acid oxidizes the hydrogen halide 
and much free iodine and hydrogen sulphide are formed: 

H 2 S0 4 + 8HI -> H 2 S t + 4H 2 + 4I 2 1 . 

The action affords a rough test for an iodide (c/. pp. 3, 200). 

Powdered sodium iodide and concentrated phosphoric acid (c/. 
p. 196), when warmed, give pure hydrogen iodide very slowly. 

The best method is one similar to that described under hydrogen 
bromide. Phosphorus and iodine unite directly to form PI 3 . 
This is a yellow solid which is violently hydrolyzed by water and 
gives phosphorous acid and hydrogen iodide: 

PI 8 + 3H 2 -> P(OH) 3 + 3HI t • 

If excess of water, which dissolves hydrogen iodide, is avoided, the 
latter goes off in a continuous stream in a gaseous condition. The 
apparatus shown in Fig. 76 may be used. The mixture of iodine 
and red phosphorus is placed in the flask and the water in the 
funnel. 

Still another method of making hydrogen iodide is frequently 
employed when a solution of the gas in water is required, and not 
the gas itself. Powdered iodine is suspended in water, and hydro- 
gen sulphide gas (q.v.) is introduced through a tube in a continuous 
stream. The iodine dissolves slowly in the water, I 2 (solid) «=* I 2 
(dslvd), and acts upon the hydrogen sulphide, which likewise dis- 



202 COLLEGE CHEMISTRY 

solves, H 2 S (gas)<=^H 2 S (dslvd). Sulphur separates in a fine 
powder, S (dslvd) <=± S (solid), and hydrogen iodide is formed in 
accordance with the equation: 

H 2 S + I 2 -»2HI + S|. 

This action takes place, however, only in presence of water, al- 
though the water does not appear in the equation. The solution 
is freed from the deposit of sulphur by filtration, and may be con- 
centrated to 57 per cent of hydriodic acid by distilling off the water. 

Physical Properties. — Hydrogen iodide is a colorless gas with 
a sharp odor. Its molecular weight is 128, and it is therefore much 
heavier than air, the average weight of whose molecules is 28.955 
(p. 101). It is a nonconductor of electricity, both in the gaseous 
and in the liquefied conditions. It is exceedingly soluble in water, 
so that at 10° 30 grams of water will absorb 70 grams of the gas, 
giving a 70 per cent solution (425 vols. : 1 aq). The behavior of 
this solution is similar to that of hydrogen chloride and hydrogen 
bromide (c/. p. 145). The mixture of constant boiling-point dis- 
tils over at 127° (at 760 mm.), and contains 57 per cent of hydro- 
gen iodide. 

Chemical Properties. — Hydrogen iodide is the least stable of 
the hydrogen halides. When heated it begins visibly to decompose 
into its constituents at 180°. On account of the ease with which 
it parts with the hydrogen which it contains, it can be burned in 
oxygen gas, 4HI + 2 — > 2H 2 + 2I 2 . When the gas is mixed 
with chlorine, a violent chemical change, accompanied by a flash of 
light, occurs, the iodine is set free, and hydrogen chloride is pro- 
duced, Cl 2 + 2HI — > 2HC1 + I 2 . Bromine vapor will similarly 
displace the iodine from hydrogen iodide. 

Chemical Properties of Hydriodic Acid HI, Aq. — In most 
respects the aqueous solution behaves exactly like hydrochloric 
and hydrobromic acids. With oxidizing agents, for example, such 
as manganese dioxide, it gives free iodine, just as the others (p. 158) 
give free chlorine and bromine, respectively. Here, however, the 
oxidation is so much more easily carried out, that it is slowly 
effected by atmospheric oxygen, so that hydriodic acid left exposed 
to the air gradually becomes brown (0 2 + 4HI — > 2H 2 + 2I 2 ). 



HYDROGEN IODIDE 203 

Although the dry gas is not an acid, the solution has all the ordi- 
nary properties of this class of substances (cf. p. 52). The hydro- 
gen may be displaced by metals like zinc and magnesium (p. 60). 
The acid interacts with oxides and hydroxides, forming iodides and 
water (p. 146). 

The Direct Union of Hydrogen and Iodine. — The union of 
hydrogen and iodine, giving hydrogen iodide, is a reversible re- 
action: 

2HI +± H 2 + I 2 . 

That is to say, whether we charge a tube with hydrogen iodide, 
or with an equal amount of the elements in the correct proportions 
by weight, if we place both tubes in a bath, and keep them thus at 
the same temperature, the contents of the tubes will after a time 
be identical (p. 177). At 283°, there will be 82 per cent of the com- 
pound, and 18 per cent of the uncombined elements. At 508° the 
proportions will be 76 per cent and 24 per cent, respectively. 

The proportion of the elements increases with rise in tempera- 
ture because the dissociation absorbs heat (p. 189). 

At any one temperature, say 283°, the equilibrium point can 
be displaced in either direction (p. 181). If we introduce some 
additional hydrogen (or iodine), without enlarging the tube, thus 
increasing the concentration of the hydrogen (or iodine), more 
than 82 per cent of the compound is formed. If, instead, we let 
one end of the tube project, and cool this end, the iodine con- 
denses to solid form, while the other two substances remain 
gaseous. This lowers the .concentration of the iodine in the 
gaseous mixture, and lowers the speed and force of the union of 
the elements. It does not affect the tendency to dissociation of the 
compound molecules, but, since it interferes with the formation 
of more of them, it enables the dissociation to proceed to practical 
completion. The condensation of the iodine is essentially like a 
precipitation (pp. 144, 186). 

This reaction illustrates very clearly the way in which the prog- 
ress of a reversible, chemical action is controlled by mechanical 
causes. It shows also why we do not prepare the compound by 
uniting the elements: (1) Since the elements interact as gases, 
very bulky apparatus would be required to prepare any consider- 



204 



COLLEGE CHEMISTRY 



able quantity; (2) the union is very slow, taking many hours at 
283°; (3) it is incomplete, at best, and we obtain a mixture, and 
not a pure substance. 

Note that, removing one product is, in general, more effective 
than increasing the concentration of one of the interacting sub- 
stances. The concentration of one product can be reduced to 
zero. To achieve the same effect by adding an interacting sub- 
stance, the concentration of the latter would have to be raised to 
infinity, which is impossible. 

Fluorine F 2 . 

The discussion of this element should logically have preceded 
that of chlorine, since it is, of all the members of the halogen family, 
the most active. Chlorine was taken up first, however, because 
its compounds are more familiar. Fluorine is found in nature 

chiefly in the mineral fluorite, calcium 
fluoride CaF 2 and in cryolite, a double 
fluoride of aluminium and sodium 3NaF, 
A1F 3 . 

Preparation. — When a solution of 
hydrofluoric acid is heated with man- 
ganese dioxide, oxidation does not occur 
and free fluorine is not produced. Until 
recently all efforts to isolate the element 
failed. It was perfectly understood that 
the reason of these failures lay in the 
greater chemical activity of fluorine, which 
made it more difficult of separation from 
any state of combination than the other 
halogens. Its preparation was finally 
achieved by Moissan (1886) by the de- 
composition of anhydrous hydrogen fluo- 
ride, which is liquid below 19°, by means of electricity. The 
apparatus (Fig. 77) is made of copper, which, after receiving a thin 
coating of the fluoride, is not further affected. To reduce the 
tendency to chemical union, the whole is immersed in a bath giving 
a temperature of — 23°. The electrodes are made of an alloy of 
platinum and iridium, which is the only material that can resist 




Fig. 77. 



HYDROGEN FLUORIDE 205 

the action of the fluorine. Hydrogen fluoride, like other hydrogen 
halides, is a nonconductor of electricity, and a small quantity of 
potassium-hydrogen fluoride KHF 2 has to be added to enable the 
current of electricity to pass. The fluorine is set free at the posi- 
tive electrode, and hydrogen appears at the negative. The U-tube 
is closed, after the introduction of the hydrogen fluoride, by means 
of blocks made of calcium fluoride, which is naturally unable 
further to enter into combination with fluorine. For the reception 
and examination of the fluorine gas, other copper tubes can be 
screwed on to the side neck of the apparatus, and, when necessary, 
small windows of calcium fluoride can be provided. 

Physical Properties. — Fluorine is a gas whose color is like 
that of chlorine, but somewhat paler. Its density (38) shows that 
the molecule is diatomic (F 2 ). The gas is the most diflicult of the 
halogens to liquefy. The liquid boils at — 186°. 

Chemical Properties. — Fluorine unites with every element, 
with the exception of oxygen, chlorine, nitrogen, and the members 
of the helium family, and in many cases does so with such vigor 
that the union begins spontaneously without the assistance of 
external heat. Dry platinum and gold are the elements least 
affected. It explodes with hydrogen at the ordinary temperature, 
without the assistance of sunlight. On the introduction of a drop 
of water into a tube of fluorine, the oxygen of the water (vapor) 
is instantly displaced by fluorine, and the vessel is filled with the 
deep-blue gas, ozone: 3F 2 + 3H 2 -> 3H 2 F 2 + 3 . 

Fluorine displaces the chlorine in hydrogen chloride as easily as 
chlorine in turn displaces bromine or iodine. 

Hydrogen Fluoride H 2 F 2 

Preparation. — Pure, dry hydrogen fluoride is best made by 
heating potassium-hydrogen fluoride, 2KHF 2 *± K 2 F 2 + H 2 F 2 1 . 
For ordinary purposes, however, the preparation of an aqueous 
solution is the ultimate object. Usually powdered calcium fluoride 
is treated with concentrated sulphuric acid, and the mixture dis- 
tilled in a retort of platinum or lead : 

CaF 2 + H 2 S0 4 <=± CaS0 4 + H 2 F 2 1 . 



206 COLLEGE CHEMISTRY 

The hydrofluoric acid passes over and is caught in distilled water. 
The aqueous solution thus obtained has to be kept in vessels made 
of lead, rubber, or paraffin, as glass interacts with the acid with 
great rapidity (see below). 

Physical Properties. — Hydrogen fluoride is a colorless liquid, 
boiling at 19.4°. It mixes freely with water and, on distillation, an 
acid of constant boiling-point (120° at 760 mm.) containing 35 per 
cent of hydrogen fluoride is obtained. The weight of 22.4 liters 
of the vapor varies from 20 g. at 90° and above, to 51 g. at 26°. 
At 90°, therefore, the formula is HF and at 26° probably a mixture 
of H 2 F 2 (40) and H 3 F 3 (60). Since HF is the only form which per- 
sists through a range of temperature, we say this substance shows 
association at lower temperatures. Water is spoken of as an 
associated liquid — the vapor being pure H 2 0, but the liquid a 
mixture of this along with (H 2 0) 2 and (H 2 0) 3 (p. 138). 

Chemical Properties of Hydrofluoric Acid H 2 F 2 , Aq. — 

Metals like zinc and magnesium interact with hydrofluoric acid 
with evolution of hydrogen (p. 60) . The action is less violent than 
with other halogen acids. The acid interacts with oxides and 
hydroxides, forming fluorides (p. 146). The chief difference in this 
respect which it exhibits, when compared with the other halogen 
acids, is one which leads us to assign to it the formula, H 2 F 2 . We 
may displace either one or both the hydrogen atoms in the molecule 
with a metal. Thus, one of the commonest salts of hydrofluoric 
acid is potassium-hydrogen fluoride, or the acid fluoride of potas- 
sium KHF 2 , mentioned above. In this respect the acid resembles 
sulphuric acid and other acids containing more than one replace- 
able hydrogen unit. 

The most remarkable property of hydrofluoric acid depends on 
the great tendency which fluorine has to unite with silicon, forming 
the gaseous silicon tetrafluoride. Glass (q.v.) is essentially a mix- 
ture of silicates of calcium and sodium, with excess of silica (sand) 
Si0 2 , and is rapidly decomposed by hydrofluoric acid: 

CaSi0 3 + 3H 2 F 2 -> SiF 4 t + CaF 2 + 3H 2 0, 
Si0 2 + 2H 2 F 2 -> SiF 4 | + 2H 2 0. 

In all other silicates, fluorine is substituted (p. 162) for oxygen 



THE HALOGENS AS A FAMILY 



207 



according to the same plan. The silicon tetrafluoride SiF 4 is a gas. 
The fluorides of calcium and sodium are solid and crumble away or 
dissolve. Thus the glass is completely disintegrated. The vapor 
of hydrofluoric acid, generated in the way described above from 
calcium fluoride in a lead dish, is used for etching glass. The sur- 
face of the glass is covered with paraffin to protect it from the action 
of the vapor, and with a sharp instrument portions of this paraffin 
are removed where the etching effect is desired. The vapor gives 
a rough surface where it encounters the glass (test for a fluoride). 
In this way, the graduation on thermometers, burettes, and other 
pieces of apparatus, is marked. The aqueous solution makes 
smooth depressions on the surface of glass. It is used for removing 
sand from metal castings and for cleaning the exteriors of buildings 
of granite and sandstone. 



The Halogens as a Family 

The most noticeable fact is that, if we arrange the halogens in 
order in respect to any one property, chemical or physical, the other 
properties will be found to place them in the same order. In the 

table, the sixth column contains the weight of the element dissolv- 
ing in 100 c.c. of water (15°). The last column, cal. KX, gives the 
heat of formation of one gram-molecule of the potassium halide. 



Element. 


Atomic 
Weight. 


State. 


Boiling- 
point. 


Color. 


Solubility. 


Cal. KX. 


Fluorine . . 
Chlorine . . 
Bromine . . 
Iodine . . . 


19.0 

35.5 

79.9 

126.9 


gas 
gas 
liquid 
solid 


-187° 

- 34° 

+ 59° 

184° 


yellow 
yellow 
brown 
violet 


"7\2" 
3.2 
0.015 


118,100 

104,300 

95,100 

80,100 



It will be seen that, as the atomic weight increases, the boiling 
point (b.-p.) rises, the color deepens, the solubility diminishes, and 
the heat of union with potassium becomes smaller. The vigor 
with which the halogens unite with hydrogen and the metals is 
greatest with fluorine and diminishes progressively until we reach 
iodine. We shall see later that the affinity for oxygen, on the 
other hand, increases as we pass from fluorine to iodine. 
Although showing different degrees of activity, the halogens are 



208 COLLEGE CHEMISTRY 

closely alike in chemical nature. That is, the relations (p. 163) 
they show when in combination are similar. When united with 
hydrogen and the metals, they are all univalent. In their oxygen 
compounds, however, they exhibit a higher valence. Their oxides 
interact with water to give acids, and they are therefore non- 
metals (p. 94). They are strongly electro-negative (pp. 55, 194), 
as non-metals all are. Their hydrides, when dissolved in water, 
are all active acids. This, and their valence, distinguish the 
halogen family from other groups of non-metals. Thus, oxygen 
and sulphur are bivalent (and the latter sexivalent also), and the 
hydrides of oxygen (H2O and H 2 2 ) and of sulphur (H 2 S) are very 
feeble acids. 

Order of Activity of the Non- Metals. — The way in which 
chlorine displaces bromine and iodine from bromides (p. 194) and 
iodides (p. 199), and bromine, in turn, displaces iodine suggests an 
order of activity for non-metals. It was noted that oxygen dis- 
places iodine from hydriodic acid (p. 202) and that iodine displaces 
sulphur from hydrogen sulphide (and all other sulphides). The 
order is, therefore, F, CI, Br, O, I, S. 

Compounds of the Halogens with Each Other 

Iodine unites directly with chlorine to form two compounds. 
The more familiar one is a red crystalline substance, iodine mono- 
chloride IC1. Another compound, IC1 3 , is made by the use of 
excess of chlorine. Iodine unites with bromine to form the com- 
pound IBr, while a compound with fluorine, IF5 is supposed to 
exist. None of these compounds are particularly stable, and some 
of them decompose easily. 

Exercises. — 1. What impurities is commercial iodine likely to 
contain? In what way does heating with potassium iodide (p. 199) 
free it from these? 

2. Classify all the chemical actions in this chapter according as 
they belong to one or other of the ten kinds (p. 166). 

3. What are the relative volumes of the gases in the interaction 
of chlorine with hydrogen bromide (p. 198), and hydrogen iodide 
(p. 202), respectively? 



THE HALOGENS AS A FAMILY 209 

4. Tabulate, more fully and specifically than is done in the sec- 
tion on "The Halogens as a Family," (a) the physical properties, 
(b) the chemical properties, (c) the chemical relations, of the mem- 
bers of this group. 

5. Construct the equation on p. 199 by the use of partial 
equations as in the example on p. 195. 

6. What are the relative volumes of fluorine and ozone in the 
action of the former upon water (p. 205)? 

7. What relative volumes of chlorine and iodine vapor must be 
taken to make the two chlorides of iodine (p. 208), respectively? 

8. At a given temperature, would increasing the pressure in a 
mixture of hydrogen and bromine vapor render the union more or 
less complete? Is the action more complete at a high or at a low 
temperature? 



CHAPTER XVI 
DISSOCIATION IN SOLUTION 

The employment of interacting substances in the form of solu- 
tions is so constant in chemistry, and the reasons for this are so 
cogent, that we must now resume the discussion of this subject 
(cf. p. 121). 

The present chapter will be devoted to giving the proofs that 
the molecules of acids, bases, and salts, in aqueous solutions, are 
actually dissociated into parts by the solvent. This will be shown 
by consideration, successively, of certain peculiarities in the 
chemical behavior, in the freezing-points and in the boiling-points 
of the solutions of these substances. We shall see that these parts 
coincide in composition with the radicals. 

Some Characteristic Properties of Acids, Bases, and Salts, 
Shown in Aqueous Solution. — Acids all contain hydrogen 
(p. 53). In aqueous solution, if soluble, they are sour in 
taste, they turn blue litmus red, and their hydrogen is displaced 
by certain metals (p. 53), and has the properties of a radical. 
By the last statement is meant that it very readily exchanges 
places with other radicals in reversible double decompositions (p. 
147) . Amongst the acids mentioned have been : hydrochloric acid 
HC1, sulphuric acid H2SO4, hypochlorous acid HCIO, acetic acid 
HCO2CH3. Many other bodies, like sugar, kerosene, and alcohol, 
contain hydrogen also, but not one of them shows all of these 
properties. 

Again, all salts are made up of two radicals, and the reversible 
double decompositions into which they enter with acids, bases, 
and other salts, consist in exchanges of these radicals. Other 
substances may include the same combinations of atoms, but in 
their actions these groupings are often disregarded. Thus, sodium 
chloride NaCl and silver nitrate AgN0 3 exchange radicals com- 
pletely (p. 147) and, in dilute solution, hydrogen chloride and 

210 



DISSOCIATION IN SOLUTION 211 

sodium-hydrogen sulphate do so partially (p. 143). But sodium 
chloride and nitroglycerine C3H 5 (N0 3 )3 do not interact at all. The 
latter is not a salt, although it contains the same proportion of 
nitrogen to oxygen as does any nitrate. 

All bases contain hydroxjd OH as a radical, combined with some 
positive radical. Potassium hydroxide KOH is soluble and active, 
zinc hydroxide Zn(OH) 2 and many others, however, are insoluble. 
Bases all exchange radicals readily in double decomposition with 
salts and acids. Other substances, like alcohol C 2 H 5 OH, may 
contain hydroxy! , but do not interact readily with salts like NaCl, 
and are not bases. 

The Influence of Water and Other Solvents. — It is chiefly 
in aqueous solution that these special properties of acids, bases, and 
salts become apparent. Their behavior is often quite different in 
the absence of this solvent. If, for example, we mix together dry 
ammonium carbonate (NH 4 ) 2 C03 and partially dehydrated, solid 
cupric nitrate Cu(N03) 2 , and apply heat, a violent interaction 
begins. An immense cloud of smoke and gas is thrown out of the 
tube, and the substance remaining is either black, or reddish, in 
parts, according to the proportions of the substances employed. 
The residue contains cupric oxide, and sometimes red cuprous 
oxide Cu 2 0. The gas is tinged red by the presence of nitrogen 
tetroxide N0 2 , while a more careful examination would show that 
it contained carbon dioxide, nitrogen, nitrous oxide N 2 0, water 
vapor, and perhaps still other products. 

The contrast, when these substances are dissolved in water before 
being brought in contact with one another, is very great. A pale- 
green precipitate is formed at once, and rapidly settles out. On 
examination, this turns out to be a carbonate of copper (basic), 
while evaporation of the solution furnishes us with ammonium 
nitrate. There are only two main products, and the essential 
part of the action in solution may be represented by the equation : 

(NHO2CO3 + Cu(N0 3 ) 2 ->CuC0 3 1 + 2NH4NO3. 

In the interaction between the dry substances the molecules are 
completely disintegrated, the whole change is very complex, and 
it takes a good deal of time. In the action in water no heating is 
required, the substances are neatly broken apart, certain groups 



212 COLLEGE CHEMISTRY 

of atoms, which we call radicals, are transferred as wholes from 
one state of combination to another, and the rearrangement takes 
place instantaneously in a machine-like manner. Contrasts like 
this between the interactions of anhydrous and dissolved bodies 
are very common. 

Many compounds, however, do not show any change in be- 
havior when dissolved in water. Sugar, for example, is, as a rule, 
more readily acted upon in the absence of any solvent. Then 
again, while water is not the only solvent which has the effect we 
have just described, the majority of solvents, if they affect chemi- 
cal change at all, simply retard it. Thus the union of iodine and 
phosphorus in the absence of a solvent takes place spontaneously 
with a violent evolution of heat. When the elements are dissolved 
in carbon bisulphide, before being mixed, the action is much milder, 
although the product is the same (phosphorus tri-iodide). The 
diminution in the concentration of the ingredients has decreased 
the speed of the action in the normal way (p. 181). That water 
and some other solvents have a specific influence tending to in- 
crease the apparent activity of certain classes of substances, shows 
that a special explanation of the phenomenon must be found. 

Summing up these points we see that the peculiarity of acids, 
bases, and salts in aqueous solution is that the action is complete 
as soon as the solutions have been mixed, and that each compound 
always splits in the same way. Thus, cupric nitrate always gives 
changes involving Cu and NO3 and never interacts so as to use 
CuN 2 and 3 , or Cu0 2 and N0 2 , as the basis of exchange. Simi- 
larly, dilute acids always offer hydrogen in exchange, and so nitric 
acid behaves as if composed of H and N0 3 , and sulphuric acid as 
if composed of 2H and SO4, and never as if made up of HSO and 
H0 3 , or H 2 S and 4 . The sour taste and the effect upon litmus 
seem to be properties of this easily separable hydrogen, for they 
are shown only by acids. The result is that we can make a list of 
the units of exchange, such as H, OH, N0 3 , C0 3 , S0 4 , Cu, K, and 
CI, employed by acids, bases, and salts in their interactions. The 
molecule of each compound of these classes contains at least two 
of them. Even when these units contain more than one atom, 
their coherence is as noticeable within this class of actions, as is 
the permanence of the atomic masses themselves in all actions. 

The question raised in our minds is whether solution in water 



DISSOCIATION IN SOLUTION 213 

alters the character of the molecule, simply by producing a sort of 
plane of cleavage in it which creates a predisposition to a uniform 
land of chemical change, or whether it actually divides the molecules 
into separate parts consisting of the above units of exchange, and 
leaves subsequent chemical actions to occur by cross-combination 
of these fragments. The fact that the dissolved substances can be 
recovered by evaporation of the liquid does not demonstrate that 
they have not been decomposed temporarily while in solution. 
The alteration which the water produces, whatever it be, will 
naturally be reversed when the water is removed. Since our 
question involves nothing but the counting of particles, the num- 
ber of which would be much greater in the event that actual sub- 
division of molecules is the explanation, it can be answered by a 
studjr of the physical properties of solutions. Several physical 
properties can be used, and they give concordant answers to the 
question. We shall confine ourselves here, however, mainly to the 
evidence furnished by the freezing-points and boiling-points of 
solutions. 

Laws of Freezing-Point Depression. — Every pure liquid 
has a definite temperature at which it freezes. Thus, pure water 
freezes at 0° and benzene at 5.48°. As we have seen (p. 134), how- 
ever, the presence of a foreign, dissolved body lowers the freezing- 
point, although the "ice" which separates usually consists of 
crystals of the pure solvent only. 

The depression in the freezing-point is directly proportional to 
the weight of dissolved substance in a given amount of the solvent. 
The depression is inversely proportional to the amount of solvent. 
Thus, if we double the concentration of the solution, the depression 
in the freezing-point is doubled. Thus, in one set of experiments, 
solutions of sugar containing 11.4 g., 22.8 g., and 34.2 g. of sugar 
to 100 g. of water were found to freeze at —0.62°, —1.24°, and 
— 1.86°, respectively. 

Further, equal numbers of molecules of different solutes in the 
same quantity of solvent give equal depressions. Or, in other words, 
the depression is proportional to the concentration of the molecules 
of the solute. Thus, solutions containing 342 g. of sugar C12H22O11,* 
or 46 g. of alcohol C 2 H 6 0, or 74 g. of methyl acetate CH3(C 2 H 3 02), 
* 12 X 12 + 22 X 1 + 11 X 16 = 342. 



214 COLLEGE CHEMISTRY 



in 1000 g. of water, being weights which contain equal numbers of 
molecules, show a depression below the freezing-point of water of 
about 1.86° in each case. That is, such solutions all freeze close 
to —1.86°. This depression, produced by a mole of the solute in 
1 1. of solvent, is called the molecular depression constant, and has 
a different value for each solvent. For solutions of the same 
molecular concentration in benzene the depression is 4.9°, in 
phenol (carbolic acid) 7.3°. Combining these facts in one ex- 
pression: 

The observed depression 1 _ 1 » fi0 Wt. of Solute 1000 

in an aqueous solution J ' Mol. Wt. of Solute Wt. of Solvent 

For other solvents, the corresponding value of the depression con- 
stant must be substituted for 1.86°. 

These laws describe the facts most exactly when the solutions 
are dilute. They hold only when there is no chemical interaction 
between solute and solvent. Even so, however, acids, bases, and 
salts dissolved in water present many apparent exceptions and must 
be discussed separately (see below). 

It will be noted that, when the other factors in the foregoing 
equation are known or observed, the molecular weight of the solute 
may be determined. The fact makes possible the determination of 
this constant for substances which are not volatile (see Hydrogen 
peroxide) . 

Abnormal Freezing -Point Depression: Dissociation in 
Solution. — The substances which present the most conspicuous 
exceptions to the above rules are acids, bases, and salts in aqueous 
solution. With most of these, the depression produced is abnormal ; 
it is greater than we should expect from the concentration of the 
solution. Thus, in an actual experiment, two equi-molar solu- 
tions were compared. One contained one mole (74 g.) of methyl 
acetate, and the other one mole (58.5 g.) of sodium chloride, each 
dissolved in 2000 g. (2 liters) of water. The freezing-points 
observed were: 

Pure water 0.000° Pure water 0.000° 

Sol. of methyl acetate . -0.970° Solution of salt .... - 1.678° 

Depression . 970° Depression 1 . 678° 

0.970° 



. 



Excess depression by salt . 708 c 



DISSOCIATION IN SOLUTION 215 

The solution of methyl acetate, as it contained only 0.5 moles of 
the solute per liter of water, showed, as it should do, about half the 
average molecular depression (1.86°, p. 214). This is typical of 
the class of substances showing normal behavior. Sugar, alcohol, 
and hundreds of other substances, in solutions of the same molar 
concentration, would have given the same value. 

The freezing-point of the salt solution, however, was much lower. 
If this solution had really contained the same concentration of dis- 
solved molecules as the other solution, its depression would like- 
wise have been 0.970°. The number of molecules in the solution 
must therefore have been greater than we should have expected 
from the number of molecules taken. In other words, a portion 
of the molecules of the salt must have been broken up, and the 
excess depression, 0.708°, must have been due to the extra mole- 
cules produced by dissociation. Now sodium chloride molecules 
cannot give more than two particles each, and the depression is 
proportional to the number of particles. It follows, therefore, 
that Iff, or 0.732 (73.2 per cent) of the molecules were dissociated: 

(27 per cent) NaCl<=± (Na) + (CI) (73 per cent). 

This result is typical also. Acids, bases, and salts, of which one 
mole is dissolved in two liters of water, are found to give irregular 
values, all more or less in excess of 0.970°. Those which contain 
but two radicals, like sodium chloride NaCl and potassium nitrate 
KX0 3 , give values between 0.970° and 2 X 0.970°. Substances 
like calcium chloride Ca(Cl) 2 and sodium sulphate (Na) 2 S04 give 
depressions approaching three times the normal value: their 
molecules contain three radicals. The excess depression depends, 
therefore, upon the number of particles which each molecule can 
furnish, and upon the proportion of all the molecules which is 
dissociated into these fragments. 

In the case of an acid, base, or salt, the depression is not strictly 
proportional to the concentration. Thus, one mole of salt in four 
liters of water does not give half the depression of the two-liter 
solution (1.678° -f- 2 = 0.839°) but somewhat more (about 0.844°). 
The same method of calculation indicates, therefore, a greater 
degree of dissociation (about 79 per cent) in the more dilute solu- 
tion (see Ionic equilibrium). 

Acids, bases, and salts, so far as they are soluble in materials like 



216 COLLEGE CHEMISTRY 

toluene, benzene, chloroform, and carbon bisulphide, exhibit 
simply normal depressions in these solvents. It appears, there- 
fore, that, in many solvents, dissociation does not take place. In 
common experience it is encountered only in solutions in water, 
and in alcohol. 

Abnormal Boiling 'Point Elevation. — We have seen (p. 135) 
that 342 g. of sugar, or an equal number of molecules of glycerine 
C 3 H 8 3 (92 g.), dissolved in 1000 c.c. of water, will elevate the 
boiling point from 100° to 100.52°. One molecular weight of 
sodium chloride (58.5 g.), however, will elevate the boiling-point 
of the water 0.87° instead of 0.52°. The effect is 0.35°, or 67 per 
cent greater, indicating dissociation of this proportion of the NaCl 
molecules. In more dilute solutions, the elevation is relatively 
greater. Salts containing more than two radicals, like Ca(Cl) 2 , 
give elevations of more than twice the normal value. In solvents 
like benzene and carbon disulphide, however, no abnormal eleva- 
tion is observed with any solute. The phenomena are, in fact, 
parallel with those connected with the freezing-point. 

Other Evidence of Dissociation. — The freezing-point and 
boiling-point are only two oifour properties of solutions which can 
be used for determining the numbers of molecules present. Nu- 
merous measurements show that aqueous solutions of acids, bases, 
and salts have also abnormal osmotic pressures (c/. p. 135). The 
electrical conductivity is the fourth property which gives the 
required information (see Chap. XVIII). Now, when we observe 
the behavior of the same solution in each of these four ways, and 
calculate the degree of dissociation from the result of each measure- 
ment, we find that the values obtained are usually identical, within 
the limits of error to which the methods are liable. Thus the in- 
dications of dissociation found in the chemical behavior of acids, 
bases, and salts (pp. 211-213) are fully confirmed by a study of the 
physical properties of their solutions. 

Applications: The Constitution of Solutions of Acids, 
Bases,, and Salts. — The composition of solutions which are nor- 
mal or abnormal, in respect to osmotic pressure, freezing-point, and 
boiling-point, may be shown thus: 



DISSOCIATION IN SOLUTION 217 



Solutes. 


Dissolved in 

Water, Alcohol, 

etc. 


Dissolved in 
Toluene, Chlo- 
roform, etc. 


Acids, bases, salts 


Abnormal 
Normal 


Normal 


Other substances . . 


Normal 







It appears that water and some other solvents have the power of 
decomposing acids, bases, and salts. Such solvents have, in fact, 
an effect on these materials that resembles, outwardly at least, the 
effect which heat has on many substances (e.g., p. 117), they cause 
dissociation: CaCl 2? ±(Ca) + 2(C1). 

In consequence of this, our view of the nature of an aqueous solu- 
tion of hydrogen chloride HC1, or common salt NaCl, or sodium 
hydroxide NaOH, or any of the substances of the classes which 
these represent, may now be stated in definite terms. Such a solu- 
tion contains, besides undivided molecules of the solute, at least 
two other kinds of material, H, Na,* CI, OH, etc., which result from 
the breaking up of the molecules. We shall see that these sub- 
divisions of the original molecules have distinct physical and chemi- 
cal properties of their own. The descriptions of the " properties" 
of the solutions, as they used to be given in chemistry, were really 
a confused statement of the properties of the different components 
of a mixture. 

The free radicals, of whose existence we have thus become con- 
vinced, constitute a new set of materials (with appropriate names. 
See p. 236). Thus the hydrogen radical of acids, although a form 
of uncombined hydrogen, differs totally from the gas which is com- 
posed of the same material. The gas has no sour taste or effect 
upon litmus; these are properties of the free radical. The gas is 
very slightly soluble in water, while the hydrogen radical exists as a 
separate substance only in solution. Again, substances with the 
composition of the radicals NO3 and SO4 are not known at all 
except in solutions. 

Exercises. — 1. What depression in the f.-p. of water will be 
produced by dissolving 10 g. of bromine in 1 kg. of this solvent? 

* The objection that separate atoms of sodium could not remain free in 
water, will be disposed of later. 



218 COLLEGE CHEMISTRY 

2. What depressions in the f .-p. of benzene and of phenol would 
be ^produced by 10 g. of bromine to 1 kg. of the solvent, if no 
chemical action took place? 

3. What is the molecular depression-constant of a solvent in 
which 5 g. of iodine in 500 g. of the solvent lowers the f.-p. 0.7°? 

4. What is the degree of dissociation of zinc sulphate, if 5 g. of it 
dissolved in 125 g. of water produce a lowering of 0.603° in the f.-p.? 

5. In a decinormal solution, potassium chloride is 86 per cent 
ionized. What is the freezing point of this solution? 



CHAPTER XVII 

OZONE AND HYDROGEN PEROXIDE 

A fresh, penetrating odor, resembling that of very dilute 
chlorine, was noticed by van Marum (1785) near an electrical 
machine in operation. Schonbein (1840) showed that the odor 
was that of a distinct substance, which he named ozone (Gk., to 
smell), and he discovered a number of ways of obtaining it. It is 
very questionable whether there is any ozone in the air, excepting 
temporarily in the immediate neighborhood of a natural or artificial 
discharge of electricity. 

Preparation of Ozone 3 . — The most satisfactory way of 
preparing ozone is to allow electric waves to pass through oxygen. 
The apparatus (Fig. 78) consists of two co-axial glass tubes, be- 
tween which the oxygen flows. The waves are generated by con- 




necting an outer layer of tinfoil on the outer tube, and an inner 
layer of tinfoil in the inner tube with the poles of an induction coil. 
With dry, cold oxygen, about 7.5 per cent of the gas is turned into 
ozone. 

Ozone is found in the oxygen generated by electrolysis of dilute 
sulphuric acid (p. 55). It arises during the slow oxidation of 
phosphorus by the air, resulting, probably, from the decomposition 
of unstable, highly oxidized bodies which are formed during the 
action. Oxygen containing 15 per cent of it is produced by the 
interaction of fluorine and water (p. 205). Ozone is formed also 

219 



220 COLLEGE CHEMISTRY 

when a jet of burning hydrogen, or an electrically heated loop of 
platinum wire is immersed in liquid oxygen. This method shows 
that ozone is formed at high temperatures, and survives when 
cooled suddenly by the liquid oxygen. 

Physical Properties of Ozone, — Ozone is a gas of blue color. 
It boils at —119°, so that when a mixture of oxygen and ozone is 
led through a U-tube immersed in liquid oxygen ( — 182.5°), the 
ozone collects in the tube as a deep-blue fluid. Ozone is much 
more soluble in water than is oxygen. At 12°, 100 volumes of 
water would dissolve 50 volumes of the gas at one atmosphere 
pressure. 

Chemical Properties of Ozone. — The density of ozone is 
one-half greater than that of oxygen. Its molecular weight is 
therefore 48, and its formula 3 . Being formed with absorption 
of energy, ozone is most stable at very high temperatures (Van't 
Hoff's law, p. 188). 

30 2 + 61,400 cal. *± 20 3 . 

When produced in cold oxygen, by energy from electric waves, it 
decomposes slowly. But this change, like all others, is hastened 
by raising the temperature. Equilibrium, with almost no ozone, 
is reached instantly at 250-300°. Liquid ozone sometimes de- 
composes explosively. As the equation shows, three volumes of 
oxygen give two of ozone. 

Ozone is a much more active oxidizing agent than oxygen. Mer- 
cury and silver, which are not affected by the latter, are converted 
into oxides by the former. Silver gives the peroxide, Ag 2 2 , thus: 

2Ag + 20 3 -> Ag 2 2 + 20 2 . 

Paper dipped in starch emulsion containing a little potassium 
iodide is used as a test for ozone: 

3 + 2KI + H 2 -> 2 + 2KOH + I 2 . 

The iodine gives a deep-blue color to the starch (c/. p. 200). This 
test, however, will not distinguish ozone from chlorine or hydrogen 
peroxide, and may, therefore, be used only in the absence of these 
substances, 






OZONE AND HYDROGEN PEROXIDE 221 

Ozone also removes the color from many of the vegetable color- 
ing matters and artificial dyes. It should be understood that the 
great majority of the complex compounds of carbon are colorless. 
Even a slight chemical change, affecting only one or two of the 
atoms in a complex molecule, is thus almost sure to give a color- 
less or much less strongly colored material. Indigo, Ci6Hi N 2 O 2 , 
which has a deep-blue color, is an example of a vegetable dye that 
is also made artificially. When ozonized air is bubbled through a 
dilute solution of this dye (as maigo-carmine), the indigo is oxidized 
to isatin C 8 H 5 N0 2 , and the color disappears (see below). 

Ozone is used commercially in bleaching oils, waxes, ivory, 
flour, and starch. It is employed also for sterilizing drinking 
water in Petrograd, Lille, and other cities. For this purpose, how- 
ever, bleaching powder is less expensive. 

Oxidizing Agents, and Explanation of their Activity. — 

When ozone turns into oxygen much heat is liberated (equation, 
above). Ozone possesses, therefore, much more internal energy 
than does oxygen. On this account it brings to the task of oxidiz- 
ing any substance more energy than does oxygen itself, and is there- 
fore more efficient. Thus, free oxygen does not interact in the 
cold with indigo, or with silver or potassium iodide (see above), 
while ozone oxidizes them rapidly. 

The heats of reaction show the difference very clearly. In 
equation (2), 1800 cal. is the amount of heat which would be 
liberated if indigo could be oxidized to isatin by oxygen gas. 
When ozone is used, we obtain, in addition, the heat of decompo- 
sition of this substance (equation 1), so that the total heat liber- 
ated (equation 3), 63,200 cal., is 35 times as great as in equation 
(2) where free oxygen is the oxidizing agent : 

20 3 = 20 2 (+ 20) + 61,400 cal. (1) 

C16H10N2O2 + (20) = 2C 8 H 5 N0 2 + 1800 cal. (2) 

C16H10N2O2 + 20 3 = 2C 8 H 5 N0 2 + 20 2 + 63,200 cal. (3) 

By similar reasoning we explain the superiority of potassium per- 
manganate over free oxygen for oxidizing hydrochloric acid (p. 
157). Substances which are more active oxidizers than is free 
oxygen may be called active oxidizing agents. 



222 COLLEGE CHEMISTRY 

It should be noted that when ozone acts as an oxidizing agent, 
usually only one of the atoms of oxygen in each molecule plays 
this part, and oxygen gas is formed. This is illustrated in all the 
three examples cited in the preceding section. 

Allotropic Modifications. — We have seen that a substance 
may exist in more than the three regular states, solid, liquid, and 
gaseous. When a simple substance shows more than one form, 
in the same state, like oxygen and ozone, we call them allotropic 
modifications. 

Hydkogen Peroxide H 2 2 

Hydrogen peroxide is found in minute amounts in rain and snow. 
It is formed in small quantities, in a way not at present fully under- 
stood, when moist metals, like zinc, lead, and copper, rust. 

Preparation of Hydrogen Peroxide. — When sodium peroxide 
is added, a little at a time, to a cold dilute acid, hydrogen peroxide 
is set free and remains dissolved in the liquid. 

Na 2 2 + 2HC1 ±5 2NaCl + H 2 2 . 

When hydrated barium peroxide (Ba0 2 ,8H 2 0) is shaken with 
cold, dilute sulphuric acid a similar action takes place: 

Ba0 2 + H 2 S0 4 *=► BaS0 4 J, + H 2 2 . 

Phosphoric acid is largely employed instead of sulphuric acid in the 
commercial manufacture of hydrogen peroxide, and great care is 
taken to precipitate the other products and all impurities from the 
solution. 

An aqueous solution is also obtained by passing carbon dioxide 
through barium peroxide suspended in water : 

Ba0 2 + C0 2 + H 2 f=> BaC0 3 j + H 2 2 . 

Pure hydrogen peroxide is isolated from any of these solutions by 
distillation under reduced pressure. To secure the low pressure, 
the ordinary distilling apparatus (Fig. 51, p. 93) is made com- 
pletely air-tight, and is connected by a branch tube with a water- 
pump. Hydrogen peroxide is much less volatile than water, but 
decomposes into water and oxygen violently at 100°. Hence the 



HYDKOGEN PEROXIDE 223 

lower pressure is required to make possible its volatilization at a 
temperature below this point. At 26 mm. pressure, the water 
begins to pass off first (at about 27°). The last portion of the 
liquid boils at 69° and is hydrogen peroxide. 

By evaporating the commercial (3 per cent) solution at 70°, a 
liquid containing 45 per cent of hydrogen peroxide may be made 
without much loss of the material by volatilization. 

Physical Properties, — Hydrogen peroxide (100%) is a syrupy 
liquid of sp. gr. 1.5. It blisters the skin and, when diluted, has 
a disagreeable metallic taste. It has been frozen (m.-p. —2°). 

Chemical Properties, — Hydrogen peroxide (100 per cent) is 
very unstable, and decomposes slowly even at —20°. The dilute 
aqueous solution, when free from impurities, keeps fairly well. 
The presence of a trace of free acid increases its stability. Free 
alkalies and most salts assist the decomposition; hence the neces- 
sity for purifying the commercial solution. Addition of powdered 
metals, of manganese dioxide, or of charcoal (contact action) 
causes effervescence even in dilute solutions, and oxygen escapes: 

2H 2 2 -> 2H 2 + 2 . 

Since the substance cannot be vaporized, even at low pressure, 
without some decomposition, its molar weight has been determined 
by the freezing-point method. The freezing-point of a 3.3 per cent 
solution in water was —2.03°. Substitution of these data in the 
formula (p. 214) gives 31.8 g. as the molar weight. Now the for- 
mula HO corresponds to a molar weight of 17 and H 2 2 to one of 
34. It is evident, therefore, that the latter is the correct formula. 

Hydrogen peroxide, in solution in water, is a feeble acid. As an 
acid it enters into double decomposition readily, and the peroxides 
are salts with the negative radical 2 n (peroxidates) . Thus, 
when hydrogen peroxide is added to solutions of barium and 
strontium hydroxides, the hydrated peroxides appear as crystalline 
precipitates : 

Sr(OH) 2 + H 2 2 <=± 2H 2 + Sr0 2 . 

The precipitation involves another equilibrium: Sr0 2 -f- 8H 2 +± 
Sr0 2 ,8H 2 (solid). 



224 COLLEGE CHEMISTRY 

The formation of a beautiful blue substance by the action of 
hydrogen peroxide upon dichromic acid is used as a test. The 
test is carried out by adding a drop of potassium dichromate to an 
acidulated solution of the peroxide. The acid interacts with the 
dichromate, giving free dichromic acid: 

H 2 S0 4 + K 2 Cr 2 7 *± H 2 Cr 2 7 + K 2 S0 4 . 

The blue substance, which is very unstable and quickly decom- 
poses, is a perchromic acid. A blue, crystalline perchromic acid 
(HO) 4 Cr(OOH) 3 , which decomposes above —30°, has been pre- 
pared. The blue substance has the property, unusual in inor- 
ganic compounds, of dissolving much more readily in ether than 
in water. It is also much less unstable when removed from the 
foreign materials in the aqueous solution. Hence the test is 
rendered more delicate by extracting the solution with a small 
amount of ether. In the ethereal layer the color of the com- 
pound is more permanent, as well as more distinctly visible on 
account of the greater concentration. 

Hydrogen peroxide is a much more active oxidizing agent than is 
free oxygen. This would be expected from the fact, that it con- 
tains so much more internal energy than the water and oxygen 
into which it decomposes (p. 223), that 23,100 cal. are liberated in 
the decomposition of one mole. Thus, it liberates iodine from 
hydrogen iodide, an action which, in presence of starch emulsion 
(cf. p. 200), is used as a test for its presence: 

2HI + H 2 2 -> 2H 2 + I 2 . 

It converts sulphides into sulphates. The white lead (q.v.) used in 
paintings is changed by the hydrogen sulphide in the air of cities to 
black lead sulphide: Pb 3 (OH) 2 (C0 3 ) 2 + 3H 2 S -> 3PbS + 4H 2 + 
2C0 2 . This may be oxidized to white lead sulphate by means of 
hydrogen peroxide: 

PbS + 4H 2 2 -» PbS0 4 + 4H 2 0, 

and in this way the original tints of the picture may be practically 
restored. Organic coloring matters are changed into colorless sub- 
stances by an action similar to that of ozone (cf. p. 221). Hence 
hydrogen peroxide is used for bleaching silk, feathers, hair, and 
ivory, which would be destroyed by a more violent agent. The 



HYDROGEN PEROXIDE 225 

products of its decomposition, being water and oxygen only, are 
harmless, and, on this account, it is used in disinfecting (destroy- 
ing organisms in) sores, and as a throat wash. 

Hydrogen peroxide exercises the functions of a reducing agent in 
special cases, also. Thus, silver oxide is reduced by it to silver: 

AgsO + H 2 2 -> 2Ag + H 2 + 2 . 

A solution of potassium permanganate, in which the permanganic 
acid has been set free by an acid : KMJ1O4 + H 2 S04 <=^ HMnC>4 + 
KHSO4, is rapidly reduced. The permanganic acid, with excess 
of sulphuric acid, tends to undergo the first of the following changes, 
provided a substance, such as hydrogen peroxide, is 'present which 
can take possession of the oxygen that would remain as a balance: 

2HMn0 4 + 2H 2 S0 4 -> 2MnS0 4 + 3H 2 (+ 50). (1) 

(50) + 5H 2 2 -> 5H 2 + 50 2 . (2) 

2HMn0 4 + 2H 2 S0 4 + 5H 2 2 -> 2MnS0 4 + 8H 2 + 50 2 . 

Exercises. — 1. What volume of ozone will be taken up by 100 
c.c. of water at 12° from a stream of oxygen containing 7.5 per cent 
of ozone (p. 129)? 

2. At what temperature will a ten per cent aqueous solution of 
hydrogen peroxide freeze (p. 214)? 

3. Write the thermochemical equations for oxidation of indigo 
by hydrogen peroxide (pp. 221, 224). 

4. How many times its own volume of oxygen gas will a 3 per 
cent solution of hydrogen peroxide give off when treated with: 
(a) platinum powder (p. 223); (6) sulphuric acid and potassium 
permanganate? 



CHAPTER XVIII 
IONIZATION 

Introductory, — As we have seen, acids, bases, and salts, when 
dissolved in water, interact with one another by interchanging 
radicals (p. 148). We have also learned that the same solutions 
have abnormal values for their freezing-points and for two other 
properties. These facts indicate dissociation into the radicals (p. 
216). Now precisely these solutions have a property which is not 
shared by any other solutions, namely, that of being conductors of 
electricity and suffering chemical decomposition by the passage of the 
current. Such solutions are called, in consequence, electrolytes, 
and the process is named electrolysis. Now the natural inference 
from the foregoing facts is that the electricity is carried by the 
liberated radicals. Our first aim in the present chapter is to show, 
by a study of the chemical changes taking place in electrolysis, that 
this inference is correct. We then proceed to discuss the nature of 
ions as a kind of molecules. Next, we devote ourselves to the 
explanation of electrolysis, to the equilibrium between the ions and 
the remaining, undissociated molecules, and to conductivity phe- 
nomena as a means of measuring the fraction ionized. Finally, we 
deduce the relation between extent of ionization and chemical 
activity. 

Incidentally, the facts to be given provide the means of under- 
standing the electrolytic processes, many of them of great impor- 
tance in chemical industries, to which frequent reference is made in 
later chapters. 

Non-Electrolytes. — To clear the ground, we should first note 
the fact that only solutions (as a rule) possess both of the properties 
in question, namely that of conducting and that of being decom- 
posed by the current. Some substances, notably the metals and 
materials like carbon, are conductors. But they are not changed 
chemically by the current. Again, single substances, even when 
they are such as, if mixed, yield electrolytes, are not conductors at 

226 






IONIZATION 227 

ordinary temperatures. Thus hydrogen chloride, whether gaseous 
or liquefied, is a nonconductor, and water is a very feeble conduc- 
tor, although the solution of the two conducts exceedingly well. 
Dry acids, bases, and salts, except when at a high temperature and 
fused, are likewise nonconductors. Furthermore, even amongst 
solutions, not all are conductors. Solutions of sugar and other 
substances of the same class (p. 213), which have normal freezing- 
points, are nonconductors. Only solutions of acids, bases, and 
salts in certain specified solvents, of which the commonest is water, 
are electrolytes at ordinary temperatures. 

Chemical Changes Taking Place in Electrolysis: at the 
Electrodes. — When the wires from a battery are attached to 
platinum plates immersed in any electrolyte (e.g., Fig. 65, p. 155), we 
observe that the products appearing at the two electrodes are 
always different. They may be of several kinds physically, and 
will be secured for examination variously according to their nature. 
Thus, when they are gases, which are not too soluble, they may be 
collected in inverted tubes filled with the solution. Solids, if in- 
soluble in the liquid, will either remain attached to the electrode or 
fall to the bottom of the vessel as precipitates. Soluble substances, 
on the other hand, will usually not be visible. They may be 
handled by interposing a porous partition of some description 
which will restrain the diffusion of the dissolved body away from 
the neighborhood of the electrode, while not interfering appreciably 
with the passage of the current. Surrounding one electrode with 
a porous battery jar is a convenient method for effecting this. 

Of the various illustrations which we have encountered, the elec- 
trolysis of hydrochloric acid (p. 155) happens to have been the only 
one which delivered both components of the solute with a minimum 
of modification at the electrodes: 

Neg. Wire, H 2 < H.C1 >C1 2 , Pos. Wire. 

Hydrogen does not interact with water, and chlorine interacts very 
incompletely, so that the molecular substances H 2 and Cl 2 are 
promptly formed from the elements H and CI which are liberated. 
The chlorides, bromides, and iodides of those metals which do not 
interact with water (p. 60) give equally simple results: 

Neg. Wire, Cu< Cu.Br 2 >Br 2 , Pos. Wire. 



228 COLLEGE CHEMISTRY 

Thus the solute seems to be split into its radicals and, in elec- 
trolysis, the radicals, if they do not interact with water, are set 
free. A substance thus set free is called a primary product of 
the electrolysis. In the foregoing instances both products are 
primary. 

Usually the chemical change is more complex. Thus, when 
dilute sulphuric acid is electrolyzed, hydrogen and oxygen are 
liberated at the negative and positive electrodes, respectively. 
But these products do not account for the whole of the constitu- 
ents (H 2 S0 4 ). We therefore proceed to examine the materials in 
solution round the electrodes. It is found that, as the action 
progresses, sulphuric acid accumulates round the positive wire, 
while the liquid in the neighborhood of the other pole is gradually 
depleted of this substance. In view of this fact we easily explain 
the phenomenon. Evidently the substance divides into its radi- 
cals, H and S0 4 , but S0 4 , not being a known substance, must 
interact with the water to produce sulphuric acid and oxygen: 
2S0 4 + 2H 2 -> 2H2SO4 + 2 . The whole change may therefore 
be tabulated as follows: 

Neg. Wire, H 2 < H 2 .S0 4 >0 2 andH 2 S0 4 , Pos. Wire. 

Hence the hydrogen is a primary product, but the oxygen and sul- 
phuric acid are secondary products. All acids give hydrogen alone 
at the negative electrode, whatever may be the product at the 
positive. 

If we electrolyze cupric nitrate solution, we obtain a red deposit 
of metallic copper on the negative plate and at the positive end 
oxygen and nitric acid are formed. We infer, therefore, that the 
division of the original molecule was into Cu and NO3, but that the 
latter interacted with the water: 4N0 3 + 2H 2 > 4HN0 3 + 2 : 

Neg. Wire, Cu< Cu.(N0 8 ) a — >Qa and HN0 3 , Pos. Wire. 

With a solution of potassium chloride we find hydrogen and 
chlorine appearing at the negative and positive electrodes, re- 
spectively. Litmus paper, however, shows the presence in the 
solution of a base (potassium hydroxide, KOH) at the negative 
end. We infer that the parts of the parent molecules are K and 
CI. The former, since it resembles sodium in being much more 
active than hydrogen (p. 60), is more difficult to liberate. Hence 



IONIZATION 229 

hydrogen is liberated instead, and potassium hydroxide remains 
in the liquid: 2K + 2HOH -» 2KOH + H 2 : 

Neg. Wire, H 2 and KOH< K.C1 >C1 2 , Pos. Wire. 

We are confirmed in this explanation when we employ a solution 
containing a mixture of salts of copper and silver. The latter, 
being the less active metal, is first deposited, alone. The copper 
is liberated only after all the silver has been set free. 

Having now before us the results of electrolyzing some typical 
substances, we bring these results into relation with the facts 
described in Chapter XVI. Acids contain hydrogen which pos- 
sesses certain specific properties (p. 210), and by electrolysis all 
acids divide so as to give up this constituent alone at one electrode. 
The evidence that the other radical has different electrical proper- 
ties which cany it to the opposite plate is conclusive. Again, salts 
undergo double decomposition in which they exchange radicals 
with acids, bases, and other salts (p. 211), and we find that it is 
these very radicals which are withdrawn from the solution by the 
influence of the electricity. Furthermore, the radicals exist free in 
the solution, being formed by dissociation of the molecules (p. 216). 
Hence the function of the electricity seems simply to consist in sifting 
apart the two kinds of free radicals which each solution contains. 
It only remains for us to explain in detail the sifting action of the 
current. Before turning to the explanation of this phenomenon, 
however, there is one question which may be answered in passing. 
Since a solution may eventually be cleared of all the hydrochloric 
acid, for example, which it contains, we should like to know how 
the free radicals in the center of the cell reach the electrodes. 

Ionic Migration, — To know how the free radicals reach the 
electrodes, all that is necessary is to take a material, one (or both) 
of whose radicals is a colored substance, and watch the move- 
ment of the colored material as it drifts towards the electrode. 
Most salts which give colored solutions are suitable. In very 
dilute cupric sulphate solution, for example, a freezing-point 
determination shows that the depression has practically double 
the normal value. In other words, the dissociation into the 
radicals, CuS0 4 ^(Cu) + (S0 4 ), is almost complete. Now, the 
blue color of the solution cannot be due to the few remaining 



230 



COLLEGE CHEMISTKY 



molecules of C11SO4, for anhydrous cupric sulphate is colorless. 
Nor is it due to the color of the (S0 4 ) radicals, for dilute potassium 
sulphate and dilute sulphuric acid are both colorless. On the other 
hand, all cupric salts, in dilute solution, have the same tint. The 
color is therefore that of the free cupric radical (Cu). In order 
most clearly to see the motion of the cupric radical, we place the 
cupric sulphate solution in the middle of the space between the 
electrodes, and place between it and the latter a colorless con- 
ducting solution. The motion of the blue material across the 
boundary may then be easily observed. 

The most convenient arrangement is to dissolve the cupric sul- 
phate in warm water containing about 5 per cent of agar-agar (a 
gelatine obtained in China from certain sea-weeds), and to fill with 
this mixture the lower part of a U-tube (Fig. 79). The setting of 
the jelly prevents subsequent mixing of the cupric sulphate system 
-^ of materials with the rest of the 

/^1 ^\ j^ y^V filling of the tube, and the conse- 

quent disappearance of the bound- 
ary. A few grains of charcoal are 
scattered on the surface of the 
jelly to mark the present limits of 
the colored substance, and a solu- 
tion of some colorless electrolyte, 
such as potassium nitrate, is 
added on each side. To prevent 
agitation of the liquid by the 
effervescence at the electrodes, it 
is well to use agar-agar with the 
lower part of the colorless liquid 
also. The whole is finally placed 
in ice and water, to prevent melt- 
ing of the jelly by the heat caused by resistance, and the current 
is then turned on. 

After a time, we observe that the blue cupric radicals ascend 
above the mark on the negative and descend away from it on the 
positive side. In each case there is no shading off in the tint. The 
motion of the whole aggregate of colored radicals occurs in such a 
way that, if the contents of the tube were not held in place by the 
jelly, we should believe that a gradual motion of the entire blue 




Fig. 79. 



IONIZATION 231 

solution was being observed. With a current of 110 volts, and a 
16-candle-power (one-half ampere) lamp in series with the cell, the 
effect becomes apparent in a few minutes. 

Although the (S0 4 ) radicals are invisible, we may safely infer 
that they are drifting towards the positive electrode. Indeed, this 
can be demonstrated by interposing a shallow layer of jelly con- 
taining some barium salt a little distance above the charcoal layer 
on the positive side. When the (SO4) reaches this, barium sul- 
phate BaS04 begins to be precipitated and the layer becomes 
cloudy. In similar ways the progress of other colorless ions may 
be rendered visible. 

It appears, therefore, that electrolysis is not a local phenomenon, 
going on round the electrodes only, but that the whole of the 
products of the dissociation of the solute are set in motion. It is on 
account of this remarkable property of traveling or migrating 
towards one or other of the electrodes that the individual atoms 
(like Cu), or groups of atoms (like S0 4 ), have been named ions 
(Gk., going). The term was first applied by Faraday to the 
materials liberated round the electrodes. 

Different ionic substances move with different speeds when pro- 
pelled by the same current. The hydrogen radical of acids (H) is 
the most speedy, the hydroxyl radical of bases (OH) comes next. 
These are, respectively, about five and two and one-half times as 
fast as any other ions. The actual speeds of several ions, in dilute 
solutions at 18°, when driven by a potential difference of 1 volt 
between plates 1 cm. apart, expressed in cm. per hour is: H 10.8, 
OH 5.6, Cu 1.6, S0 4 1.6, K 2.05, CI 2.12. 

The Nature of Ions: Faraday's Law, — That the molecules 
of certain classes of substances, although seemingly without chemi- 
cal interaction with the water in which they are dissolved, should 
nevertheless be decomposed by the influence of the water, is 
strange, but not inconceivable. Heating produces a somewhat 
similar effect on many substances. The novel fact, for which an 
explanation is demanded, is that the molecules of the products of 
the dissociation appear to be attracted by electrically charged 
plates, which have been lowered into the solution, while molecules 
of dissolved sugar, for example, are not so attracted. Now the only 
bodies which we find to be conspicuously attracted by electrically 



232 COLLEGE CHEMISTRY 

charged objects are bodies which are already provided with electric 
charges of their own. Thus we are led to add the idea that sub- 
stances which undergo dissociation in solution divide themselves 
into a special kind of electrically charged molecules. 

Since the solution, as a whole, has itself no charge, equal quan- 
tities of positive and negative electricity must be produced: 

HC1 1=± H+ + CI" NaCl <=> Na+ + CI" NaOH *± Na+ + OBT. 

This means that bivalent radicals, on dissociation, will become ions 
carrying a double charge and trivalent ions must carry a triple 
charge : 

CuCl 2 ±± Cu++ + 2C1" CuS0 4 <± Cu++ + S0 4 = 
K 2 S0 4 <=± 2K+ + S0 4 = FeCl 3 *± Fe+++ + 3C1" 

In these equations, the coefficients multiply the charges as well as 
the radicals bearing the charges, and it will be seen that the num- 
bers of + and — charges produced by each dissociation are equal. 
Hence, univalent ions all possess equal quantities of electricity, and 
other ions bear quantities greater than this in proportion to their 
valence. This is an inevitable inference from the electrical neu- 
trality of all solutions. An ion is therefore an atom or group of 
atoms bearing an electric charge. 

This conclusion is confirmed by actual measurement. When 
hydrochloric acid is electrolyzed, 35.46 g. (= CI) of chlorine are 
liberated for every 1.008 g. (= H) of hydrogen. But when cupric 
chloride CuCl 2 is substituted, for every 35.46 g. ( = CI) of chlorine 
set free, only 31.78 g. (= \ Cu = \ 63.57) of copper is deposited. 
The law, discovered by Faraday, is that : equal quantities of electric- 
ity liberate equivalent quantities of the ions (equivalent, p. 65, 
not atomic or molecular). 

To show that this view of the nature of the ions is adequate, we 
next apply it to the explanation of the phenomena of electrolysis. 
After that some seeming objections will be discussed. 

Application to the Explanation of Electrolysis, — A bat- 
tery is a machine which maintains two points, its poles, or two 
wires connected with them, at a constant difference of potential. 
One cell of a lead storage battery, for example, maintains a poten- 
tial difference of about two volts. When the wires are joined, 



IONIZATION 



233 




directly or indirectly, the poles are immediately discharged, but the 
cell continuously reproduces the difference in potential by generat- 
ing fresh electricity. Now the effect of immersing two plates, one 
of which is kept by the battery at a definite positive potential and 
the other at a definite negative potential, into a liquid filled with 
multitudes of minute, suspended bodies, already highly charged, 
may easify be foreseen. 

The figure (Fig. 80) will convey some idea of the behavior of the 
parts of a system such as we have imagined. The electrodes are 
marked — and +. The negatively charged plate attracts all the 
positively charged particles Cathode + Anode 

in the vessel and, although __ «— cation = Ag . _j_ 

these particles were in con- 
tinuous and irregular mo- 
tion, they at once begin to 
drift toward the plate in 
question. On the other 
hand, the negatively 
charged particles are re- 
pelled by this plate and 
attracted by the positive 

plate, so that they drift in the opposite direction. Those which are 
nearest each plate, on coming in contact with it, will have their 
charges of electricity neutralized by the opposite charge on the plate, 
turning thereby into the ordinary free forms of the matter of which 
they are composed. The continuous removal of the electrical charges 
of the plates through contact with ions of the opposite charge fur- 
nishes occasion for recharging of the plate from the battery, and thus 
gives rise to a continuous current in each wire. Again, the continu- 
ous drifting of positively and negatively charged particles in oppo- 
site directions through the liquid, constitutes what, in the view of 
all external means of observation, appears to be an electical current 
in the liquid also. A magnetized needle, for example, which is de- 
flected when brought near to one of the wires of the battery, is in- 
fluenced in the same way by being brought over the liquid between 
the electrodes. The illusion, so to speak, of an electric current is 
complete, although in reality it is a convection of electricity that is 
taking place. Furthermore, the quantity of electricity being trans- 
ported across any section of the whole system is the same as that 



Fig. 80. 



234 COLLEGE CHEMISTRY 

across any other, whether this section be taken through one of the 
wires, through the electrolyte, or even through the battery at any 
point. As fast as the ions are thus annihilated as such, the undis- 
sociated molecules (mingled with the ions, but not shown in the fig- 
ure) dissociate and produce fresh ones, as in all chemical equilibria. 
Eventually, by continuing the process long enough, if the substances 
set free are actually deposited and do not go into solution again 
in any form, the liquid can be entirely deprived of the whole of the 
solute which it contains. 

The analogy to the transportation of a fluid like water is notice- 
able, although not complete. Water may be transported in three 
ways. It. may flow through a pipe, it may pass by pouring freely 
from one container to another, and it may be carried in vessels. 
Thus a stream of water, essentially continuous, might be arranged, 
in which part of the passage took place through the pipes, part by 
pouring from the pipes into buckets, and part by the carrying of 
those buckets between the ends of the pipes. The quantity of 
water passing a given point per minute in this system would be the 
same at every part, although the actual method by which the water 
was transported past the various points might be different. In 
such a disjointed circuit we suppose the electricity to move when 
carried from a battery through an electrolytic cell. It flows in 
the wire, passes by discharge between the pole and the ion, and is 
transported upon the ions in the liquid. The parallel is imperfect, 
however, because we have used the conception of two electric fluids 
and because the ions are already charged in the solution, and before 
any connection with the battery is made. They do not, so to speak, 
transport the electricity of the battery, but their own. 

Questions Suggested by this Explanation. — 1. The ques- 
tion was raised (p. 217), as to how we can imagine separate atoms 
of sodium to exist in water without acting upon it, as the metal 
sodium usually does. But the ions of sodium in sodium chloride 
solution are not metallic sodium. They bear large charges of 
electricity. They possess an entirely different, and in fact, by 
measurement, much smaller amount of chemical energy than free 
sodium. And, as we have seen, the properties of a substance are 
determined as much by the energy it contains as by the kind of 
matter. Metallic sodium and ionic sodium are, simply, different 



IONIZATION 235 

substances. Besides, when metallic sodium acts on water, it turns 
into the ionic sodium of sodium hydroxide (Na + + OH~^± NaOH). 
Ionic sodium (Na + ) from sodium chloride is, therefore, already in 
the very state which metallic sodium reaches by interaction with 
water, and is in no need of trying to enter that state. 

2. We think of hydrogen chloride and common salt as exceed- 
ingly stable substances, and are averse to believing that precisely 
these compounds should be highly dissociated by mere solution in 
water. But it must be remembered that in solution they undergo 
chemical change very easily, and it is only in the dry form that 
they show unusual stability. 

3. Again, why do not the ions combine, in response to the at- 
tractions of their charges? The answer is that they do combine, 
but the rate at which combination takes place is no greater than 
that at which the molecules decompose, so that on the whole the 
proportion of ions to molecules remains unchanged. 

4. It might appear that the idea that bodies could retain high 
charges in the midst of water is contrary to all experience. It 
must be remembered, however, that the molecular, pure water, 
which separates the ions from one another, is a perfect nonconduc- 
tor. The moisture which covers electrical apparatus and causes 
leakage of static electricity is not pure water, but a dilute solution, 
containing carbonic acid (p. 91) and materials from the glass of 
which the apparatus is made (p. 92). It conducts away the 
charge electrolytically, by means of the ions it contains, and not 
by itself acting as a conductor. 

5. Finally, when we dissolve an electrically neutral salt in water, 
whence do the radicals obtain the electric charges? We now know 
that an atom, say, of sodium, contains a minute nucleus of positive 
electricity, which contains most of the mass of the atom. Outside 
of this nucleus, there are particles of negative electricity, called 
electrons (q.v.), each having a mass about one-eighteen hun- 
dredth (tsW) of that of an atom of hydrogen. An ion of chlorine 
(Cl~) consists, therefore, of an atom of chlorine plus one electron 
(CI + e). An ion of sodium is an atom of sodium minus one 
electron (Na — e) and has thus an excess of one unit positive 
charge in the nucleus. When these two ions combine, the result- 
ing molecule NaCl is neutral. 



236 



COLLEGE CHEMISTRY 



Resume and Nomenclature. — The dissociation of molecules 
into ions is named ionization. The substances of the three classes 
which alone are ionized may be designated ionogens. An ion may 
be defined as, a molecule bearing negative or positive charges of 
electricity in proportion to its valence, and formed through the 
dissociation of an ionogen by a solvent like water. 

Each molecule of the solute gives two kinds of ions with opposite 
charges. These two are forthwith distinct and independent sub- 
stances, save that the attractions of the charges prevent any con- 
siderable separation by diffusion. They differ from non-ionic 
substances of the same material composition when such are known. 
The electrical charge is one of the essential constituents and, when 
it is removed, the properties alter entirely. Thus we have two 
kinds of hydrogen, gaseous molecular hydrogen (H 2 ), and ionic 
hydrogen (H+), with entirely different chemical properties. 

The radicals and their chemical behavior are real, and all the 
peculiarities of aqueous solutions of acids, bases, and salts are 
experimental facts. We now have experimental knowledge of 
the minute parts of bodies. Molecules are units which are not 
commonly disintegrated by vaporization (p. 102); ions, those 
which are not commonly disintegrated in double decomposition in 
solution; atoms, those which are not commonly disintegrated in 
any chemical action. The ionic explanation was first suggested 
as an hypothesis by Svante Arrhenius, a Swedish chemist, in 
1887. 

Since ionic hydrogen, ionic chlorine, etc., are entirely different in 
physical and chemical properties from the corresponding free ele- 
ments, they should receive separate names. When it is incon- 
venient to say " ionic hydrogen," " ionic nitrate radical" (NO3""), 
etc., the following names will be used for the ionic substances: 



Sym- 
bol. 


Name of 
Substance 


Anion of 


Symbol. 


Name of 
Substance. 


Cation of Salts of 


SO4- 

cr 

Hsor 

OH" 


Sulphate-ion 
Chloride-ion 
Hydrosulphate-ion 
Hydroxide-ion 


Sulphates 
Chlorides 
Bisulphates 
Hydroxides 
(bases) 


Na+ 
F e-H-+ 

NH 4 + 
Fe++ 
H+ 


Sodium-ion 

Ferric-ion 

Ammonium-ion 

Ferrous-ion 

Hydrogen-ion 


Sodium 
Ferric iron 
Ammonium 
Ferrous iron 
Hydrogen (acids) 



IONIZATION 237 

In using these terms, note that sodium-ion (with the hyphen) is the 
name of the substance, and not of the charged atom. When 
speaking in terms of ions as particles, therefore, we may not say 
"a sodium-ion," any more than we should say "an ionic sodium" 
or "ionic sodiums." To describe the charged molecule, we must 
write "a sodium ion," "sodium ions," "chlorate ions," etc. 

Faraday distinguished the two kinds of material which proceed 
with and against the positive current by name. His terminology is 
still used. Ions which proceed in the same direction as the positive 
current (Fig. 80) are called cations (Gk., down). Such are H + , 
CU++, K + , NH4 4 ". They are metallic elements, or groups which play 
the part of a metal. The electrode (Gk., a path for electricity) upon 
which they are deposited, the negative electrode, is spoken of as 
the cathode (Gk., the way down). 

The particles which move in the direction of the negative current, 
and against that of the positive, are named anions (Gk., up). The 
ions Cl~, N0 3 - , S0 4 = , Mn0 4 ~ are of this kind. They are usually 
composed of non-metals, although sometimes, as in Mn04~~, the con- 
stituents may be partially metallic. They are set free at the posi- 
tive electrode, which is therefore named the anode (Gk., the way 
up). Chemists speak of metals and non-metals as positive and 
negative elements, respectively, even when electrical relations are 
not directly in question, and ions are not concerned. 

Actual Quantities of Electricity Concerned, — The units 
of electrical energy are the coulomb, which is the unit of quantity, 
and the volt, which is the unit of difference of potential (or pressure, 
so to speak). Faraday's law has to do only with the former. 
Equal numbers of coulombs liberate equivalent weights of all 
elements, but different voltages are required to decompose differ- 
ent compounds, according to their stability (see Chap. XXXIX). 

To liberate 1.008 g. of hydrogen, or one equivalent of any other 
element, 96,540 coulombs of electricity are needed. The charges 
on 1.008 g. of hydrogen ions must, therefore, equal this amount. 
There are 6.07 X 10 23 molecules of hydrogen in 22.4 liters (H 2 ) and 
therefore in 2.016 g. of the gas. A simple calculation shows there- 
fore that each coulomb is distributed over about 63 X 10 17 ions 
of hydrogen. 

A current of 1 coulomb per second is called 1 ampere. Thus, 



238 COLLEGE CHEMISTRY 

the current passing through a 1-amp. lamp (or 2 half-ampere 16- 
c.p. lamps in parallel) will liberate 1.008 g. (11.2 liters) of hydro- 
gen in 96,540 seconds, or 26 hours and 49 minutes. The same 
current will liberate 107.88 g. of silver (Ag 1 ), or 31.78 g. of copper 
(Cu n /2) from cupric sulphate in the same time. A current of 
5 amperes will accomplish the same result in one-fifth of the time. 

Applications: Ionic Equilibrium. — Since the ions are chemi- 
cally different from their parent molecules, their formation repre- 
sents a variety of chemical change. The change does not involve 
any chemical interaction with the water. It is simply a dissocia- 
tion, i.e., reversible decomposition of the dissolved substance. 

From the fact that the proportion of molecules ionized is shown 
to become greater as more and more of the solvent is added (p. 215), 
and that removal of the solvent diminishes the proportion of ions to 
molecules, and finally leaves the substance entirely restored to 
the molecular condition, we know that this is a reversible action and 
therefore a true dissociation. The molecules and their ions adjust 
themselves like the constituents in any case of chemical equilibrium 
(pp. 177-182): 

NaCl*±Na+ + Cr. 

The chemical behavior of substances in ionic equilibrium will be 
discussed in the next chapter (see p. 249). 

* The mode of formulation previously used (p. 183) may be 
employed here. If [NaCl], [Na+], and [Cl~] stand for the molec- 
ular concentrations (numbers of moles per liter) at equilibrium of 
the molecules, and the two ions, respectively, we have an equilib- 
rium constant (cf. p. 184), in this case called the ionization 
constant: 

[Na+]X[Cr] 



K = 



[NaCl] 



When we dissolve a single substance which gives only two ions, the 
molecular concentrations of the ions are necessarily equal. When 
some other ionogen with a common ion is present, however, the 
values of [Na + ] and [Cl~] will be different. 

* The content of this paragraph is referred to in Chap. XX, but is not em- 
ployed systematically until Chap. XXXV is reached. 



IONIZATION 



239 



Applications: To the Interpretation of Conductivity 
Measurements. — We have seen that when the solution of an 
ionogen is diluted, the proportion of ions to undissociated molecules 
increases, while removal of a part of the solvent has the opposite 
effect (p. 215). Now, a change in the number of ions naturally 
modifies the capacity of the liquid for carrying electricity, so that 
observation of the changes in the conductivity of a solution, when 
the concentration is altered, supplies the simplest means of studying 
the phenomena of ionization. 

A glass trough and amperemeter * (Fig. 81) may be used to illus- 
trate this principle. The electrodes are long strips of copper foil, 
which pass down at the ends of the trough. After placing the two 




Fig. 81. 

instruments in circuit with a source of electricity, we first pour very 
pure water into the cell. With this arrangement, the ampere- 
meter does not indicate the passage of any current of electricity. 
Concentrated (36 per cent) hydrochloric acid is next cautiously 
added through a long-stemmed dropping funnel, so that it forms a 
shallow layer below the water, and the funnel is withdrawn. The 
situation at this stage is that a definite amount of hydrogen 
chloride dissolved in a small amount of water fills what was before 
a nonconducting gap in the electric circuit. The deflection of the 
needle in the amperemeter indicates that a certain current of 
electricity is able to pass through this acid. When we now stir the 

* An amperemeter of low resistance, 0.5-1 ohm, must be used. 



240 COLLEGE CHEMISTRY 

surface of the acid very gently with a thin glass rod, the ampere- 
meter instantly responds, showing an increase in conductivity. As 
we stir, the conductivity increases, and the increase ceases only 
when the liquid has become homogeneous. Introduction of an 
additional supply of water will improve the conductivity still more, 
but the effect becomes less and less, until no change on further 
dilution is perceptible. Reasoning about these effects, we perceive 
that the amount of hydrochloric acid has not altered during the 
experiment. Yet the quantity of conducting material between the 
electrodes must have become greater, for the carrying power of 
the whole has improved. We were therefore observing the progress 
of a chemical change of the nonconducting hydrogen chloride into 
conducting materials. Hydrogen chloride molecules do not carry 
electricity (p. 145), but the hydrogen and the chloride ions, into 
which it was gradually altered by chemical change during the 
stirring, do carry electricity. Furthermore, the change practically 
ceased at great dilution, for the dissociation into ions was then 
practically complete. If we could conveniently have started with 
only liquefied, dry hydrogen chloride in the cell, we should have 
observed the whole range of changes from zero to the maximum. 

When a saturated solution of cupric chloride is used instead of 
hydrochloric acid, dilution is accompanied by a similar improve- 
ment in conductivity. Here we notice, besides, that the yellowish- 
green liquid, with which we start, changes to a pale blue, as the 
yellowish-brown molecules of cupric chloride are dissociated and 
the color of the solution becomes more exclusively that of the 
copper ions. When the solution has become perfectly blue, further 
dilution is seen to affect the conductivity but slightly. 

Reasoning still further about these phenomena we see that, if we 
start with a fixed amount of a given substance, the conductivities 
at different stages of the dilution must be proportional to the numbers 
of ions, and the maximum conductivity attainable by great dilu- 
tion must represent the effect when the whole material has become 
ionic. Thus, if the conductivity at the maximum is represented, 
say, by 5, then at the dilution where the conductivity is 2, the 
proportion of the whole which is ionized is 2/5. When the con- 
ductivity becomes 4, 4/5 of the molecules are dissociated and the 
degree of ionization is 0.8. When the conductivity becomes 5, 
5/5, or all, of the molecules are dissociated. For example, in 



IONIZATION 



241 



hydrochloric acid, if we take the normal solution (p. 124) containing 
36.5 g. of acid per liter as the unit of concentration, the fractions 
ionized at various concentrations are as follows: ION, 0.17; N, 
0.78; N/10, 0.91; iV/100, 0.96. Thus, measurements of con- 
ductivity enable us to study the ionic decomposition of all ionogens, 
and to state accurately the fraction ionized, at each concentration, 
in solutions of every ionogen. This information is obviously most 
valuable, for it places us in a position to know the exact constitu- 
tion of every solution we use in the laboratory. In the following 
section the data on which such knowledge can be based is given. 
In the next chapter the mode of applying the data is explained. 

Constitution of Solutions of Ionogens: Fractions Ionized. 

— The dilute acids used in the laboratory are generally of six times 
normal (6JV) concentration. But, often, we add only a drop or 
two to a large bulk of liquid, so that the acids are commonly very 
dilute as actually employed. The solutions of salts are of different 
strengths, but the great majority are of normal (N), or even smaller 
concentrations. In practice they, also, are still further consider- 
ably diluted before use. If, therefore, we give the fractions ionized 
(total molecules of ionogen = 1) in decinormal solutions (except 
where otherwise specified), the reader will be able to estimate 
roughly the proportion of each kind of ions in any application of 
the reagent. In the case of acids containing more than one dis- 
placeable hydrogen unit, the kind of ionization on which the figure 
is based is indicated by a period. Thus H.HCO3 means that the 
whole of the ionization is assumed to be into H + and HC03~. 

FRACTION IONIZED IN 0.1JV SOLUTIONS AT 18° 
Acids 

Nitric acid 0.92 

Nitric acid (cone, 62%) . . 0.09 

Hydrochloric acid . 92 

Hydrochloric acid (cone, 

35% 0.13 

Sulphuric acid, H.H.SO4 . . 0.61 
Sulphuric acid (cone, 95%). 0.01 

Hydrofluoric acid 0.15 

Oxalic acid, H.HC 2 4 . . .0.50 
Tartaric acid, H.HT . . . .0.08 

Acetic acid (N) 0.004 

Acetic acid 0.013 



Carbonic acid, H.HCO3 . 


. 0.0017 


Carbonic acid (N/25) . . 


. 0.0021 


Hydrogen sulphide, H.HS 


. 0.0007 


Boric acid, H.H 2 B0 3 . . . 


. 0.0001 


Hydrocyanic acid .... 


. 0.0001 


Permanganic acid (iV/2) . 


. 0.93 


Hydriodic acid (N/2) . . 


. 0.90 


Hydrobromic acid (N/2) . 


. 0.90 


Perchloric acid (N/2) . . 


. 0.88 


Chloric acid (N/2) . . . 


. 0.88 


Phosphoric acid, H.H2PO4 


. 0.27 


Water 


. 0.0*1 



242 



COLLEGE CHEMISTRY 



Bases 



Potassium hydroxide . . .0.91 

Sodium hydroxide 0.91 

Barium hydroxide . 77 

Lithium hydroxide (AT) . . 0.63 
Tetramethylammonium hy- 
droxide (N/16) 0.96 



Potassium chloride . . . .0.86 

Potassium nitrate . 83 

Potassium acetate . 83 

Potassium sulphate . . . .0.72 
Potassium carbonate . . . .(0.71) 

Potassium chlorate . 83 

Ammonium chloride . . . .0.85 
Sodium chloride (N) . . . .0.66 
Sodium chloride (N/2) . . .0.74 

Sodium chloride . 84 

Sodium nitrate 0.83 

Sodium acetate . 79 

Sodium sulphate . 70 



Ammonium hydroxide . . . 0.013 
Strontium hydroxide (iV/64) . 93 
Barium hydroxide (iV/64) . . 92 
Calcium hydroxide (AT/64) .0.90 
Silver hydroxide (2V/1783) .0.39 
Water 0.0 6 1 



Salts 

Sodium bicarbonate, 

Na.HC0 3 0.78 

Sodium phosphate, Na2.HP04 . 73 

Sodium tartrate . 69 

Barium chloride 0.77 

Calcium sulphate (N/100) .0.64 

Cupric sulphate . 39 

Silver nitrate . 81 

Zinc sulphate 0.40 

Zinc chloride 0.73 

Mercuric chloride ... (<0.01) 
Mercuric cyanide .... Minute 



In addition to their use in showing the nature of the reagents 
employed in the laboratory (p. 241), these numbers show also to 
what extent any pair of ionic substances will unite when mixed (see 
pp. 247, 251), and they likewise indicate the chemical activity of 
the ionogens when in solution (see next section). 



Relation of Ionization to Chemical Activity. — These 
tables may be used for reference. The import of the following 
general statements, drawn from the tables, should be memorized: 

1. Salts, with the exception of those of mercury, are all well 
ionized. In actions involving their ions, salts are therefore all of 
the same order of activity, for a dilute solution of every salt contains 
a large amount of the ionic components. 

2. Acids show the most extreme differences in their degrees of 
ionization. That is to say their solutions must contain very differ- 
ent concentrations of hydrogen-ion. Since their activity as acids 
depends on this substance (p. 217), and the activity of a substance 
is proportional to its concentration (p. 182), it follows that acids 
will show very great differences in apparent chemical activity. At 



IONIZATION 243 

this point, therefore, we emerge from semi-physical discussion of 
the subject and reach something of definite, practical application 
in chemical work. 

The data show that acids may be divided roughly into four 
classes with different degrees of acidic activity: 

(a) The ionization in decinormal solution exceeds 70 per cent; 
e.g., nitric acid and hydrochloric acid. These are the acids which 
are chemically most active, for their solutions contain a relatively 
high concentration of hydrogen-ion. 

(b) The ionization is between 70 and 10 per cent; e.g., sulphuric 
acid and phosphoric acid. These acids are noticeably less active, 
for their solutions contain a lower concentration of hydrogen-ion. 

(c) The ionization is between 10 and 1 per cent; e.g., acetic acid. 
These are the weaker acids, for their solutions contain a very small 
concentration of hydrogen-ion. 

(d) The ionization is less than 1 per cent; e.g., carbonic and 
boric acids. These are the feeble acids, for their solutions contain 
only a minute concentration of hydrogen-ion. 

3. The bases show two classes: 

(a) Ionization high; e.g., potassium hydroxide. These bases 
are active, for their solutions contain a high concentration of 
hydroxide-ion. 

(6) Ionization less than 2 per cent; e.g., ammonium hydroxide. 
These bases are weak on account of the low concentration of 
hydroxide-ion. 

4. Water is less ionized than any other substance in the list. It 
shows therefore, as we already know, usually little or no interaction 
with acids, bases, or salts, and hence is valuable as a solvent for 
these substances. Its ions are H + and OH~, and it is thus as much 
(or as little) an acid as a base. 

Exercises. — 1. With solutions of the following substances, 
state, (a) what will be the products of electrolysis, (b) whether each 
is primary or secondary, and (c) how they may be isolated in each 
case : Potassium chlorate, potassium iodide, potassium iodate, sil- 
ver sulphate, sodium peroxide. 

2. Make equations (p. 232) showing the ionic and molecular 
materials in solutions of potassium bromide, potassium bromate, 
sodium periodate, aluminium chloride, zinc sulphate. Mark the 



244 COLLEGE CHEMISTRY 

charges on the ions and give the name of each ionic substance (p. 
236). 

3. Prepare lists of other anions and cations which have been 
encountered, giving the formula and number of charges of elec- 
tricity in each case. 

4. If the conductivity of sodium chloride solution at the maxi- 
mum is 110, and at greater concentrations is as follows: N, 74.7; 
N/10, 92.5; iV/100, 103, calculate the fraction ionized at each 
concentration. 

5. If the conductivity of acetic acid solution at the maximum is 
352, and at greater concentrations is as follows: lOiV, 0.05; N, 
1.32; iV/10, 4.6; iV/100, 14.3, calculate the fraction ionized at each 
concentration. 

6. If 1 c.c. of dilute hydrochloric acid (6iV) is added to 30 c.c. 
of an aqueous solution, what is the reacting concentration of the 
acid? 

7. Classify all the acids in the table (p. 241) according to the 
four classes (p. 243). 



CHAPTER XIX 

IONIC SUBSTANCES AND THEIR INTERACTIONS 

In this chapter, after enumerating the various classes of ionqgens, 
and the various kinds of ionic substances, we discuss the interactions 
of the latter. We consider first the relations of the ionic and the 
molecular substances (in equilibrium) when a single ionogen is 
present, and then take up the ways in which such an ionic equi- 
librium is displaced. Finally, we discuss some of the useful ionic 
interactions, in which the equilibria are displaced so far that 
practically complete interaction occurs: namely, precipitation, 
neutralization, and displacement. 

The Classes of lonogens. — Acids are classified according 
to the number of hydrogen units in their molecules. Thus chloric 
acid HCIO3 is a monobasic acid, sulphuric acid H2SO4 a dibasic acid, 
and phosphoric acid H3PO4 a tribasic acid. These terms relate to 
the fact that, in neutralization (see p. 254) the acids interact with 
one, two, or three molecules of a base like sodium hydroxide. 

Bases are named in a similar way: sodium hydroxide NaOH is a 
monoacid base, calcium hydroxide Ca(OH) 2 is a diacid base. 

Salts like KC1 and Na 2 C0 3 are neutral (see acid salts, below) or 
normal salts, and NaKC0 3 and Ca(OCl)Cl (bleaching powder) are 
mixed salts. 

The most interesting classes of mixed salts are the acid salts 
(p. 206) and the basic salts. In acid salts, like NaHS0 4 (p. 141) and 
KH2PO4 (p. 196), all the hydrogen of the acid has not been replaced 
by a metal. In basic salts, like Ca(OH)Cl, part of the basic 
hydroxyl remains. 

There are also many double salts, like ferrous-ammonium sul- 
phate (NH4) 2 S04,FeS04,6H 2 0, and alum (see index), some of which 
are in common use. 

All these substances are ionogens (p. 236). The mixed and 
double salts are, naturally, dissociated into more than two ionic 
substances. 

245 



246 COLLEGE CHEMISTRY 

Ionic Substances Furnished by Acids, — The mode of 
naming ionic substances has already been given (p. 236). 

Acids, e.g., HC1, H 2 S0 4 , when dissolved in water, all furnish 
hydrogen-ion H + and a negative ionic substance (anion), e.g., Cl~, 
S04 = . The solutions differ from those of salts in the constant pres- 
ence of hydrogen-ion, and in the absence of any other positive ion. 

Hydrogen-ion H+ is a colorless substance. It is sour in taste, 
and its presence is recognized by the fact that it turns blue litmus 
red (see Indicators, below). These properties serve as tests for 
acids, as they are not interfered with by other ionic substances 
which may be present. Hydrogen-ion is univalent and, when 
combined with negative radicals of salts, gives the (molecular) 
acids. The activity of acids depends upon the concentration of 
the hydrogen-ion they furnish (p. 242), and therefore upon their 
solubility and the degree of ionization of the dissolved molecules. 
Some furnish so little hydrogen-ion that their action on litmus 
can hardly be detected. 

Ionic Substances Furnished by Bases. — Bases, e.g., KOH, 
NH4OH, Zn(OH) 2 , all furnish hydroxide-ion OH~ and some positive 
ionic substance (cation), K+, NH4+, Zn ++ . Their solutions differ 
from those of salts in the constant presence of hydroxide-ion and 
in the absence of any other anion. The more active bases, that is, 
those which are soluble and highly dissociated, so that they give a 
high concentration of hydroxide-ion, are called alkalies. Such are 
potassium and sodium hydroxides. They are often named caustic 
alkalies and, individually, caustic potash and caustic soda. The 
solutions are called lyes. 

Hydroxide-ion OH~ is a colorless substance. Properties which 
serve as tests for bases are that hydroxide-ion possesses a soapy 
taste and feeling and turns red litmus blue (see Indicators, below) . 
It is univalent, and combines with positive radicals to form 
(molecular) bases. 

Ionic Substances Furnished by Salts, — Salts furnish 
positive and negative ionic substances, which may be either simple 
or composite, Na.Cl, Na.N0 3 , NH4.CI, NH4.NO3. Some ionic 
substances are colored, CU++ (cupric-ion) blue, Cr _H ~ f ' reddish- 
violet, Co 4-1- pink, MnC>4~ (permanganate-ion) purple, Cr207 = (di- 



IONIC SUBSTANCES AND THEIR INTERACTIONS 247 

chromate-ion) orange, but most of them are colorless, K + , Na + , 
Zn 4-1 ", Cl~, I~ N0 3 ~ SO4-. They vary in taste, some being salt, 
some astringent, some bitter. The ionic materials characteristic 
of salts do not affect litmus, and individual tests are required for 
each. Usually we add some other ionic substance, with which the 
ion thought to be present combines to form an insoluble, molecular 
substance of known color, or appearance, and examine the precipi- 
tate if any appears. Thus, when the presence of chloride-ion Cl~ 
is suspected, we may add a solution containing silver-ion Ag+, 
expecting to obtain a precipitate of silver chloride AgCl (Cl~ + 
Ag + — ■> AgClJ,). I n dilute solutions of salts, the ions are almost 
always numerous in comparison with the molecules (p. 242), so 
that salts are practically all active and their solutions almost 
always respond readily to the tests for the ions they contain. 
The art of detecting the various ionic substances present in a 
solution constitutes a large part of the branch of chemistry called 
qualitative analysis. 

All the known ionic substances are found in solutions of salts. 
The only ions which are not characteristic of salts, although some- 
times occurring in their solutions (see acid and basic salts, above) , 
are hydrogen-ion H + , and hydroxide-ion OH~~. 

It will assist the reader if the following facts are kept in mind. 
The elements which can form a simple positive ion are the 
metallic elements (p. 94, and see Chaps. XXII and XXXIII). 
Non-metallic elements, like nitrogen, may be present in a positive 
ion, as in NH4 + , but never exclusively. In other words, we know 
no such substances as nitrogen sulphate, or carbon nitrate. Con- 
versely, the metals are frequently found in the negative ion, but 
never constitute it exclusively. They are then usually associated 
with oxygen, as in Mn04~, and Cr 2 07 = . 

The Ionic Equilibrium with a Single Ionogen. — In the 

ionization of a molecular substance, the chemical change is incom- 
plete and the system reaches a condition of equilibrium (p. 238). 
The action is, therefore, reversible, and there are thus two routes to 
the same equilibrium point. This fact must not be forgotten, for 
we have to consider the union of ionic substances even more often 
than the converse change. Now, the degrees of ionization of various 
ionogens tell us the location of the equilibrium point, and therefore 



248 COLLEGE CHEMISTRY 

the extent of the chemical change involved in reaching this point by 
either route, that is, either by the dissociation of molecules or by the 
union of ions. In a class of interactions, of which all are incom- 
plete, and only those are interesting and useful which approach 
completeness, we require some means of knowing which are com- 
plete and why they are so. The table of fractions ionized (p. 241) 
supplies most of the required information. 

To illustrate, take the case of a single ionogen. When we place 
hydrogen chloride in decinormal solution, 0.92 of the molecules dis- 
sociate. Conversely, when we start with the hydrogen-ion and 
chloride-ion, say by mixing two solutions, each of which contains 
one of them (along with another ion), then 1 — 0.92, or only 0.08 
of these ionic substances will combine. 

This exemplifies the case of an active acid. The following equa- 
tions show the data for six typical substances in N/10 solution, 
namely, two acids, two bases, and two salts: 

(8%)HC1 <=±H+ +Cl-(92%), (98.7%)HC 2 H 3 2 <=±H+ + C 2 H 3 2 -(1.3%) 
(9%)KOH^K+ +OH-(91%), (98.7%)NH40H^NH 4 + + OH-(1.3%) 
(16%)NaCl <± Na+ + Cl"(84%), (61 %)CuS0 4 *± Cu++ + S0 4 =(39%) 

These samples are chosen to illustrate, in each pair, the extremes. 
Thus, when potassium-ion and hydroxide-ion are brought together 
little union takes place, while with ammonium-ion and hydroxide- 
ion the union is practically complete. In the case of the soluble 
salts, however, there are almost (p. 242) no cases of considerable 
union of the ions in dilute solutions. The case of water, on the 
other hand, is one of the most extreme: 

(99.9 5 %) H 2 <± H+ + OH" (0.0 4 1%). 

Hydroxide-ion and hydrogen-ion thus unite almost completely. 

Similar reasoning enables us to handle the more complex, but 
very common case of the mixing of two ionogens. The degrees of 
ionization tell us the exact condition of each system separately, 
before mixing. The result of the mixing is best understood by 
viewing the change as consisting in a displacement of each of the 
equilibria by the action of the components of the other. We con- 
sider, therefore, next, the displacement of ionic equilibria. 

The Displacement of Ionic Equilibria, — Equilibria are 
displaced by changes which favor or disfavor one of the opposed 



IONIC SUBSTANCES AND THEIR INTERACTIONS 249 

actions (p. 180). There may be either, (1) a physical change in the 
conditions, or a chemical interaction which (2) adds to, or (3) 
removes one of the interacting substances. Each of these may be 
illustrated in turn. 

1. As an example of the first, we have the effect of changing the 
amount of the solvent (p. 215) . Adding more of the solvent reduces 
the concentration of the ionic materials and disfavors their union, 
so that it indirectly promotes dissociation. The larger the volume 
in which the ions are scattered, the less often will they meet, and 
the smaller the amount of combination. On the other hand, 
evaporating off a part of the solvent favors the encounters of the 
ions and promotes combination. When the solvent is at last 
entirely gone, the whole material is molecular. 

In cases where the ionic and molecular substances are all color- 
less, these changes can be followed only by a study of the freezing- 
points or other similar properties of the solutions (p. 216). But 
when the substances are of different colors, the changes can also be 
seen. Thus, cupric bromide in the solid form is a jet black, shining, 
crystalline substance. When treated with a small amount of 
water it forms a solution which is of a deep reddish-brown tint, 
giving no hint of resemblance to a solution of any cupric salt. This 
doubtless represents the color of the molecules. When more water 
is added, the deep brown gives place gradually to green, and finally 
to blue. The latter is the color of the cupric-ion (CU++), and is 
familiar in all solutions of cupric salts. The colorless nature of 
solutions of potassium and sodium bromides shows that bromide- 
ion (Br~) is without color. Hence, in the present instance it is 
invisible. We are thus watching the forward displacement of the 
equilibrium: 

CuBr 2 (brown) fc? Cu++ (blue) + 2Br~ 

If 1 g. of the solid is taken, it dissolves in about its own weight of 
water, and independent measurement shows that there is relatively 
little ionization. Hence the solution is deep brown. When 10 c.c. 
of water has been added, 70 per cent of the salt is ionized, and the 
solution is green. With 40 c.c. of water, only 19 per cent remains 
in molecular form, and the blue color of the cupric-ion entirely 
overbears the tint of the molecules. If we now remove the water 
by evaporation, all these changes are reversed. When 30 c.c. of 



250 COLLEGE CHEMISTKY 

the water has been driven off, the solution is green. As the evapo- 
ration of the remaining 10 c.c. progresses, the brown color appears. 
When the water is all gone, the black residue remains. Here we 
are observing the backward displacement of the equilibrium, 
CuBr 2 ±=> Cu++ + 2Br~ 

2. Cupric bromide may be used to illustrate also the chemical 
methods of displacing equilibria. Thus, we may show the effect of 
adding more of one of the reacting substances. If, at the green stage, 
we dissolve solid potassium bromide in the liquid (KBr<=±K++Br~), 
the increased concentration of bromide-ion causes more vigorous 
interaction of the ions, and the molecules, with their brown color, 
become prominent again. Adding cupric chloride increases the 
concentration of cupric-ion and has the same effect. In either 
case, renewed dilution with water reduces the concentrations of all 
the ions once more, the molecules become fewer, and the brown 
color is displaced by the blue for the second time. 

3. Finally, the displacement of the same equilibrium by remov- 
ing one of the interacting substances may be illustrated. Thus, if 
the chocolate-brown solution, in which molecular cupric bromide 
predominates, is shaken with pulverized lead nitrate (and filtered), 
two changes are noticed. A pale yellow precipitate of lead bromide 
appears (Pb 4 " 4 " + 2Br~ — > PbBr 2 [ ), and the brown color fades into 
green. Here the displacement is the opposite of the last. Instead 
of reinforcing one of the ions, we have reduced the concentration, 
and in fact almost entirely removed one of them, namely Br~. 
This has, naturally, stopped the interaction of the CU++ and Br~ 
which reproduces the brown, molecular CuBr 2 . Hence the disso- 
ciation of the latter has continued to exhaustion of the whole 
molecular material. 

The reader will find that the behavior of these ionic equilibria, 
and the way in which we discuss and explain it, are complete 
parallels of the behavior and explanation in the case of ordinary 
equilibria (pp. 185-187), which should now be reexamined. The 
illustrations in the present section, and particularly the third (c/. 
p. 203), should be considered until every feature is perfectly clear. 
They furnish the key to understanding the applications which fol- 
low. One fact must not escape notice, and that is that in none 
of the three instances was the forward action (the dissociation) 
in itself affected. The molecules of cupric bromide have, as we 



IONIC SUBSTANCES AND THEIR INTERACTIONS 251 

should expect, a certain tendency to decompose. No encounters 
between these molecules are required for mere decomposition. 
Hence their decomposition is not influenced by their nearness to, 
or remoteness from, one another (illustration 1), nor by the presence 
of any other kinds of molecules or ions (illustrations 2 and 3). 
The effect, whether it involved an apparent increase, or a diminu- 
tion of the dissociation, was always accomplished by altering the 
concentration of the ionic substances, and therefore the activity of the 
reverse action. 

Applications: Double Decomposition in Solution, — We 

are now prepared to consider the general case of mixing the solu- 
tions of two ionogens. 

When solutions of two ionized substances are mixed, the first reflec- 
tion which occurs to us is that each of these has been diluted by the 
water in which the other was dissolved, so that the first effect will 
be to increase the degree of ionization of both to a certain extent. 

The next consideration is, however, that we have produced a 
mixture of four ions, which must have at least some tendency to 
unite crosswise. Thus potassium chloride and sodium nitrate in 
dilute solution are very greatly ionized before mixing. The re- 
versible actions, represented by the horizontal pair of the following 
equations, have taken place extensively. But, by mixing the 
liquids, we have brought into presence of -^^ „ + , p,_ 

one another two new pairs of positive and XT AT ~ JL _ AT ~ _T , T . 
, • . tt 4. fi • NaN0 3 *=► N0 3 + Na + 

negative ions. Hence, two other reversi- ,^ ,* 

ble actions, the vertical ones, will be set TTNTn NT ri 

up and will proceed until a fresh equi- 
librium of all the ions with all four kinds of molecules has been 
reached. Thus far the description will fit any case of mixing 
solutions of two ionogens. 

Now, in this particular instance, what is the actual extent of such 
interaction as has occurred? To answer this question we require to 
know the proportion of molecules to ions in a solution of each of the 
four salts (p. 242). In decinormal solutions it is KC1, 14 : 86; 
NaN0 3 , 17 : 83; KN0 3 , 17 : 83, NaCl, 16 : 84, so that the salts 
are all equally well ionized. It is a good plan to add these pro- 
portions in the formulation. Furthermore, in a dilute mixture, 
such as we are considering, the proportions of ions are greater than 



252 COLLEGE CHEMISTRY 

these figures indicate. Hence, practically no chemical action has 
occurred. 

(>83%) (>84%) 
(14%)KC1 fc? K+ + CT (86%) 

(17%)NaN0 3 ±9 N0 3 ~ + Na+ (83%) 

IT It 

KN0 3 NaCl 

«17%) «16%) 

That this inference is correct is shown by independent evidence. 
Thus when the solutions of salts are mixed, no thermal effect is 
observable. This fact has been known since 1842 as Hess' law 
of thermoneutrality. Again, if the solutions are placed in a cell 
(Fig. 81, p. 239), so that the one forms a layer below the other, no 
change in conductivity is noticed when the solutions are stirred 
together. Hence no change in the number of ions has occurred. 

We conclude, then, that when two highly ionized substances are 
mixed, and the possible products are also highly ionized, soluble 
substances, then practically no chemical action occurs. This rule 
applies to dilute solutions of all soluble salts (p. 242) and to mixing 
salts with the highly ionized acids or bases. 

Conversely, when two ionized substances are mixed, an extensive 
chemical change does ensue in two cases, namely : 

1. When one of the possible products is an insoluble substance 
and precipitation occurs, for this removes the ions used to form the 
insoluble body. 

2. When one of the possible products, although soluble, is little 
ionized, as in neutralization, for this likewise removes the ions re- 
quired to form molecules of the product. We proceed, therefore, 
to discuss these two important classes of actions. 

Precipitation, — A typical case of precipitation occurs when 
we mix dilute solutions of silver nitrate and sodium chloride. 

(>83%) 
(16%) NaCl U Na+ + CI" (84%) 
(19%) AgN0 3 *=> NO-3 + Ag+ (81%) 

It it 

NaN0 3 * AgCl (dslvd) 

«17%) it 

AgCl (solid) 






IONIC SUBSTANCES AND THEIR INTERACTIONS 253 

Here, since the four substances are all salts, they are all highly 
ionized. If they were all soluble, then, in dilute solutions, perhaps 
5 per cent of each salt would be in molecules and the rest in ionic 
form. But the molecules of silver chloride are excessively insoluble. 
In all cases of precipitation, we look up the solubilities of the possible 
products (see Table of Solubilities inside the front cover). Here 
we find that one liter of water will dissolve only 0.0016 g. silver 
chloride (this quantity includes both ions and molecules). So 
the concentration of the AgCl (dslvd) becomes almost zero through 
precipitation. So far as it is in solution, however, being a salt 
and very dilute, it is practically all ionized. The precipitation 
displaces the equilibrium, for, the dissociation having thus ceased, 
those of the ions Ag + and Cl~ which combine are not replaced by 
others. Hence the silver-ion and chloride-ion almost disappear. 
This occurrence affects in turn the equilibria with Na + and N0 3 ~, 
so that the NaCl and AgN0 3 become completely ionized. Hence 
the concentrations of NaCl and AgN0 3 , of Ag + and Cl~~, and of 
the dissolved AgCl, all become practically zero at last. The 
system finally contains only a precipitate of molecular, solid silver 
chloride and a solution of the three substances, Na + + N0 3 ~ *=> 
NaX0 3 , in equilibrium. By far the greater part of this material 
in solution is the ionic, namely the Na+ and the N0 3 ~. 

To avoid a misconception, note that the answer to the question, 
"Is silver chloride a highly ionized substance?" is "Yes." Since 
it is a salt, we expect this. True, very little of it dissolves, so that 
it cannot give many ions to a solution. But little or much ionized 
refers to the proportion ionized of the material which has dissolved. 
With undissolved material ionization has nothing to do. 

It should be noted that, when the solutions are mixed, as in the 
foregoing example, strictly speaking, the chief interaction taking 
place is the production of the insoluble body. The largest part of 
the chemical action may be formulated thus: 

Ag+- + CT -> AgCl. 

The chief change that has as yet befallen the ions of sodium nitrate 
is that they have been transferred from two separate vessels into 
one. Potentially the salt has been formed. But the actual union 
of its ions, to give the second product in the molecular condition, 

Na + + NOr -» NaN0 3 , 



254 COLLEGE CHEMISTRY 

comes about only when, at some subsequent time, if at all, the 
water is evaporated away. 

The foregoing formulation and explanation apply to every case 
of mixing ionogens where precipitation occurs, that is, where the 
products are insoluble acids, bases, or salts. 

Neutralization. — We may now consider the case of mixing 
solutions of two ionogens where one is an acid and one a base. 

(>87<7) ( W) The 3 eneral P lan of a11 k" 

(8%) HC1 1, CI" °+ H+ (92%) teractions <***<* and bases 

(9%) NaOH^Na* + OH" (91%) ^ shown "\ the J^ 1 ^ 011 - 
•* It lomzation of the hydro- 

^ p. iL chloric acid reaches 0.92 in a 

( ^ i qo/ ^ si r\(\07 ^ decinormal solution, and goes 

«16/ ) (lW/ ) farther when the acid is di- 

luted with the water of another solution. That of the sodium hy- 
droxide similarly goes beyond 0.91. Thus the substances in the 
solutions before mixing are almost entirely ionic. The crosswise 
union, H+ + OH" ±=> H 2 0, however, is all but complete, for water 
is hardly ionized at all (p. 243). The materials on whose inter- 
action with the Cl~ and Na+, respectively, the maintenance of the 
molecules HC1 and NaOH depends, being thus removed, the disso- 
ciation of the acid and base promptly brings itself to completion, 
and the left sides of the equations vanish. Practically all the 
hydrogen-ion and hydroxide-ion become water, which thenceforth 
is simply a part of the solvent. The CI" and Na + , however, if 
the solution is now 1/20 normal, unite to the extent of 0.13 only. 
If it is more dilute, this union forms a still smaller factor in the 
whole change. Practically it is negligible. Now all that has been 
said of this acid and base will apply mutatis mutandis whenever 
any active, highly ionized acid and base come together. Thus 
we may write one simple equation for all neutralizations of active 
acids and bases: 

H+ + OH" -* H 2 0, 

without omitting anything essential. 

The ions of a salt are always left over from the main action, and 
may be brought together, in turn, by evaporation: Na + -f- CI"— > 
NaCl, or the liquid may be used as a solution of the pure salt. 






IONIC SUBSTANCES AND THEIR INTERACTIONS 255 

Confirmations of this Vieiv of Neutralization. — That 
these inferences are correct is shown by many facts. The most 
conspicuous of these is the fact that, when equivalent amounts of 
active acids and bases are used, the mixture is without action either 
on red or on blue litmus. It is neutral to indicators — hence the 
term neutralization applied to the operation of mixing an acid and a 
base. Specifically, the absence of effect upon litmus demonstrates 
the absence of hydrogen-ion H + and of hydroxide-ion OH - , alike, 
in the product, and confirms the theory. 

Again, a considerable thermal effect accompanies neutralization. 
But, in the cases we are discussing, that is where active bases and 
acids are employed, the heat liberated by use of equivalent weights 
(p. 124) is always the same, namely 13,700 cal. That it is always 
the same confirms our theory, for practically the whole change is 
always the formation of 18 g. of water from the ions. 

Still again, when we place the acid and base in the cell (Fig. 81, 
p. 239), so that the one forms a layer beneath the other, and watch 
the amperemeter while we mix the solutions, a marked decrease in 
the current passing through the cell is noticed. This also confirms 
our theory, for it is our belief that one-half of the ions, namely the 
H+ and OH - , disappear as such during the action. The decrease 
is, in fact, to less than half the reading before mixing, because the 
two speediest ions have been removed. 

When less highly ionized acids or bases are used, the only differ- 
ence is that there are more of the molecular materials present, 
before the solutions are mixed. But the removal of the H+ and 
OH - ions permits the molecules of the acid and base to dissociate, 
so that the final products are water and the ions of a salt, as before. 

The foregoing formulation and explanation apply to every case of 
mixing ionogens, where a very slightly ionized substance is one of 
the products, that is, when water, or a feeble acid, or a feeble base 
(pp. 242-243) is formed. 

Acidimetry and Alkalimetry. — When, as is constantly the 
case, a chemist desires to ascertain the quantity of an acid or base 
present in a solution, he uses for the purpose the interaction just 
discussed. If, for example, the problem is to ascertain the weight 
of hydrogen chloride in each liter of a specimen of hydrochloric 
acid, this can be done by neutralizing a measured portion of this 



256 



COLLEGE CHEMISTRY 



/ 



\ 



im 




acid with a solution of an alkali of known concentration. The 
volume of the latter which is required for the purpose is observed. 
If the alkali is sodium hydroxide, the action taking place is 

HC1 + NaOH -> H 2 + NaCl. 

The volume of acid is measured out into a beaker by means of a 
pipette (Fig. 82) of fixed capacity, which is filled by suction to the 

mark on the stem. Sup- 
pose the amount to be 
25 cc. The standard 
alkali solution is placed 
in a burette (Fig. 83), 
which is filled down to 
=i — the tip of the nozzle. A 
few drops of litmus solu- 
tion are now added to 
the acid, and the alkali is 
allowed to run in slowly. 
After a time, the hy- 
droxide-ion which this 
introduces will begin to 
produce a blue color, 
close to where the 
stream enters the liquid. 
This is at first dissi- 
pated by stirring, and 
the whole remains red. 
Finally, however, a 
point is reached at 
which the entire solu- 
tion assumes a tint in- 
termediate between 
blue and red. With 
one drop less of the 



Fig. 82. 



Fig. 83. 



base, it is distinctly red. With one drop more, it would become 
distinctly blue. Litmus paper of either shade dipped in this neu- 
tral solution remains unaffected. 

By the use of a standard solution of an acid in the burette, the 
quantity of a base may be determined in the same way. 



IONIC SUBSTANCES AND THEIR INTERACTIONS 257 

Standard Solutions. — The standard solutions used in this 
work are usually normal, and contain one equivalent weight of the 
alkali or acid in one liter of the solution. For more delicate work, 
decinormal (N/10) solutions may be employed. The concentra- 
tion of such a solution is called its titer, and the operation of 
analyzing another solution by means of it, titration. The value of 
standard solutions lies in the fact that, when once the solution has 
been prepared, and the exact concentration adjusted by quantita- 
tive experiments, its use does not require any weighing, and the 
measurements of volumes can be carried out with great rapidity. 

The calculation of the result is also simple. One liter of normal 
alkali contains 17 g. of available hydroxyl, and one liter of normal 
acid, 1 g. of available hydrogen (p. 124). Equal volumes of 
normal solutions will therefore exactly neutralize one another, 18 g. 
of water being formed by interaction of a liter of each. If, for 
the neutralization of the 25 c.c. of hydrochloric acid used above, 
50 c.c. of normal alkali are required, the acid is twice-normal (2N). 
When 15 c.c. are required, the acid is ^f or $N. If the actual 
weight of the acid in the latter case has to be calculated, we remem- 
ber that there are 36.46 g. of hydrogen chloride in 1 1. of a normal 
solution, and therefore 36.46 X f X T Mo g- = 0.5467 g. in 25 c.c. 
of a solution which is f-normal. 

Methods of quantitative analysis in which standard solutions are 
employed are known as volumetric methods, and are much used by 
analysts and investigators. They occupy much less time than 
gravimetric operations, in which numerous weighings have to be 
made, and are often just as accurate. The substances, like litmus, 
by whose change of color the completeness of the action is made 
known, are called indicators. 

Indicators. — Indicators are substances which, in presence of 
certain other substances, assume a very deep color, or change 
sharply from one deep color to another. Thus, phenolphthalein 
is colorless in presence of acids (i.e., hydrogen-ion), and red (when 
dilute, pink) in presence of alkalies (i.e., hydroxide-ion). Litmus, 
again, is red with acids, and blue with alkalies. The change of 
color depends upon a chemical interaction in each case, but since 
indicators are chosen for their strong coloration, the quantity of 
the acids or base used up in changing the tint of the trace of the 



258 COLLEGE CHEMISTRY 

indicator is so small as to be negligible. The common indicators 
are: 

Phenolphthaleih C20H14O4, a colorless substance and very feeble 
acid. It is not perceptibly dissociated into its ions, 

: C20H14O4 (colorless) i=? CaoHisOr (red) + H+ 

and in neutral or acid solutions is, therefore, without visible color. 
When a base is added gradually to an acid containing some of this 
indicator, the acid is first neutralized. Then, and not till then, 
the slightest excess of hydroxide-ion unites with the trace of 
hydrogen-ion from the phenolphthalem, the above equilibrium is 
displaced forwards, and a visible amount of . the red negative ion 
is formed: 

C20H14O4 (colorless) <=> C20H13O4" (red) + H+ > 
NaOH t=> Na+ + OH" J "* M2U ' 

In this more compact formulation, we show the product (H 2 0) 
from the union of the two ions which combine, but omit the prod- 
uct from the union of Na+ and C 2 oHi 3 04~, because here (since the 
product is a salt) hardly any union occurs. 

Litmus is an extract from certain lichens, first used by Boyle. 
It contains azolitmin. One of its colors is that of the molecule, 
and the other that of the ion. 

Methyl orange (CHs^NCeHi.N : N.CelltSOsNa is a complex or- 
ganic compound which gives, in acid solution, a red, and in alka- 
line solution a yellow color. 

Congo red is the sodium salt of an acid of complex structure (see 
Dyes). In neutral or alkaline solutions it is red; with acids it 
turns blue. Paper dipped in Congo red differs from litmus paper 
in that it shows gradations in color, the blue being much more 
distinct with an active acid than with a relatively weak one like 
acetic acid (p. 241). Litmus paper is equally red with all acids 
save the very feeblest. 

Displacement: The Electromotive Series. — In the preced- 
ing sections we have dealt with cases in which ionic substances 
underwent combination or ionogens dissociated. This is one of five 
kinds of ionic chemical change. Of the remaining four, ionic dis- 



IONIC SUBSTANCES AND THEIR INTERACTIONS 259 

placement is the one * that we have most frequently encountered. 
Thus, certain metals displace hydrogen from dilute acids (p. 60) : 

Zn + H 2 S0 4 -> ZnS0 4 + H 2 . 

These interactions do not occur in the absence of water (p. 53), and 
now appear in a new light, namely, as ionic actions: 
Zn + 2H+ + S0 4 = -> Zn++ + H 2 + S0 4 = 

The molecular sulphuric acid and zinc sulphate, which are small 
in amount, are omitted because they do not, as such, take part in 
the change. On looking at the equation, we perceive that the 
sulphate-ion is also unaltered by the action, and may be left out 
likewise : 

Zn + 2H+ -> Zn++ + H 2 . 

True, hydrogen-ion cannot be used alone, *or it is always accom- 
panied by some negative radical. But the latter, like the vessel in 
which the experiment is made, is part of the necessary apparatus, 
and not an interacting substance. The change has consisted in the 
ionization of the zinc, and the transfer to it of the electric charge 
of the hydrogen-ion. In terms of electrons (p. 235), each atom 
of zinc has lost two electrons (Zn — 2e = Zn ++ ) and two ions of 
hydrogen have taken up the electrons (2H+ + 2e — ■> H 2 ) . 

These statements enable us to understand why active acids, with 
zinc, give hydrogen faster than do inactive acids (p. 54). The 
former provide a higher concentration (p. 243) of hydrogen-ion, 
that is, of the real interacting substance, than do the latter. 

A similar displacement of negative ions has been met with (pp. 
194, 199). Thus, chlorine displaces bromine from solutions con- 
taining bromide-ion. 

Cl 2 + 2Br--^2Cr + Br 2 . 

The Electromotive Series. — Displacement occurs with all 
positive ions. Thus, zinc will displace other metallic elements, 
such as iron, lead, copper, and silver, from the ionic conditions, 
when it is placed in solutions of their salts: 

Zn + Cu++ -» Zn++ + Cu. 

* The discharge of an ion and liberation of its material in electrolysis 
(pp. 55, 155, 227) is another. Attention will be called to the remaining two 
when suitable illustrations occur (see pp. 270, 504). 



260 



COLLEGE CHEMISTRY 



Here the copper appears as a red precipitate. Lead, in turn, will 
displace copper and silver, but not zinc or iron. Copper will dis- 
place silver. Thus the metals can be set down in an order, such 
that each metal displaces those following it in the list and is 
displaced by those preceding it. This list is known as the electro- 
motive series of the metals, because in electrolysis of normal solu- 
Electromotive ti° ns °^ their salts, the electromotive force of the 
current required to deposit each metal is less 
than that for the metal preceding in the list. 
For present purposes, the list shows the metals 
in the order of diminishing tendency to enter the 
ionic from the elementary condition. 

The electromotive series embodies many facts 
in the behavior of the metals, and should be kept 
in mind as furnishing a key to all actions in- 
volving solutions in which a free metal is used 
or produced. It is, in fact, identical with the 
order of activity (p. 60). 

To avoid a common misconception, it must 
be noted that the electromotive series cannot be 
used to explain the tendency of one radical to 
dislodge another in double decompositions. The 
place of an element in the E.M. series defines 
its relative activity when free, and has to do only 
with actions where one free element displaces 
(p. 55) another. The influences which deter- 
mine a double decomposition (c/. pp. 143, 186) 
are such as the insolubility of a compound. 
Thus, potassium bromide solution will slowly 
convert a precipitate of silver chloride into one 
of silver bromide: AgCl + KBr -* AgBr + KC1. 
This occurs because silver bromide is the less soluble salt. But free 
bromine never displaces chlorine from binary combination with a 
metallic element. It is free chlorine that displaces combined 
bromine. 



Series of the 
Metals. 

Potassium 

Sodium 

Barium 

Strontium 

Calcium 

Magnesium 

Aluminium 

Manganese 

Zinc 

Chromium 

Cadmium 

Iron 

Cobalt 

Nickel 

Tin 

Lead 

Hydrogen 

Copper 

Arsenic 

Bismuth 

Antimony 

Mercury 

Silver 

Palladium 

Platinum 

Gold 



Non-Ionic Modes of Forming lonogens. — While ionogens 
may always be made by the union of the proper ions, they must 
nevertheless, in the absence of the solvent, be regarded as chemical 



IONIC SUBSTANCES AND THEIR INTERACTIONS 261 

substances which may be constructed, and very frequently are 
made, out of their constituents without reference to the ionic 
plane of cleavage. Thus we have incidentally observed many 
ways in which acids, bases, and salts may be prepared, that do 
not involve a union of the constituent ions and are probably not 
ionic. 

Oxygen acids can almost all be prepared from the anhydride, 
that is, the oxide of the non-metal, which is not an ionogen, and 
water. Phosphoric acid, sulphurous acid (p. 94), hypochlorous 
acid (C1 2 + H 2 — » 2HC10) , and many other acids are so 
formed. Hydrogen fluoride, chloride, bromide, and iodide are 
all producible by union of the constituent elements. Many acids 
are formed from others when the latter are decomposed; for 
example, hydrochloric acid from hypochlorous acid (p. 161). 

Bases are formed by the union of oxides of metals with water 
(p. 94). 

The dry ways of forming salts are very numerous. Thus, many 
are produced by direct union of the elements, as in the case of chlo- 
rides (p. 146), sulphides (p. 14), and other simple salts. Many are 
made by reduction or oxidation from other salts, as potassium chlo- 
ride from potassium chlorate (p. 27), or potassium perchlorate 
(q.v.) from the latter. Often a reducing or an oxidizing agent is 
used, as in making sodium nitrite (see index) from the nitrate. 
Almost all oxygen salts can be obtained by the union of two oxides, 
as calcium carbonate (see index) from calcium oxide and carbon 
dioxide. Ammonium salts are formed by combination of am- 
monia, which is not an ionogen, with acids (p. 146). 

In manufacturing commercially important salts, methods like 
the above, as well as those involving ionic actions, are very com- 
monly used. In each case the cheapest and most easily acces- 
sible materials are chosen, and the least expensive operation is 
selected. 

Exercises. — 1. Give, for each of the following, a definition, 
i.e., concise description, in terms of experimental facts: acid (pp. 52, 
158, 210, 246), base (pp. 94, 146, 246), salt (p. 246), acid salt, 
mixed salt. 

2. Give, now, a definition of the same things (see 1), in terms of 
ions. 



262 COLLEGE CHEMISTRY 

3. Name all the ionic substances whose formulae are given on 
pp. 212, 237, and classify them into anions and cations. 

4. Give a list of the specific physical and chemical properties, 
including those that can be used as tests, of: iodide-ion, sulphate- 
ion, cupric-ion, chloride-ion. 

5. Give a list of all the colorless ionic substances you can think 
of. 

6. Using the table of fractions ionized (p. 241), prepare lists of 
the pairs of ionic substances which show the greatest, and the least 
tendency to combine, and state in each case the proportion com- 
bining in decinormal solution. 

7. In the case of the green solution of cupric bromide (p. 249), 
explain in detail (p. 181) the effect of the addition of potassium 
bromide. Formulate the action (p. 251). u 

8. In the case of the chocolate-brown, concentrated solution of 
cupric bromide (p. 249), explain in detail what would happen to 
the system: (a) if metallic zinc were to be added (p. 259); (6) if 
hydrogen sulphide gas were to be led into the solution (CuS is 
insoluble). 

9. Formulate, after the models on pp. 251 and 252, and discuss 
fully, the interaction of ferric chloride and ammonium thiocyanate 
(p. 182). 

10. What is implied by the statements, that peroxides are salts 
and that hydrogen peroxide is feebly acid (p. 223)? 

11. Formulate after the model on p. 252, and discuss fully, the 
interaction of: (a) sodium peroxide and hydrochloric acid (p. 222) ; 
(6) barium peroxide and sulphuric acid. 

12. Invent an interaction of two soluble salts in which both 
products shall be insoluble (see Table of Solubilities, inside of 
front cover) and formulate it, (p. 252). 

13. For the neutralization of 77 c.c. of a certain alkaline solution, 
25 c.c. of normal hydrochloric acid are required. What is the 
normal concentration of the alkali? If the alkali was sodium 
hydroxide, what weight of the substance was present? If the 
alkali was barium hydroxide, what weight of it was present? 

14. Formulate (p. 259) the actions of iron and of aluminium on 
dilute hydrochloric acid. 

15. Formulate (p. 259) the displacements of iodine by chlorine 
and by bromine (p. 200). 



IONIC SUBSTANCES AND THEIR INTERACTIONS 263 

16. Which metals (p. 260), besides platinum, would be most 
likely to form suitable electrodes for an electrolytic cell? 

17. To which classes of ionic actions do those of iodine on hy- 
drogen sulphide (p. 201), and of calcium on cold water (p. 50), 
belong? 



CHAPTER XX 

SULPHUR AND HYDROGEN SULPHIDE 

Occurrence. — Free sulphur is found in volcanic regions in 
Sicily, where it is mixed with gypsum and other minerals and occu- 
pies the pores of pumice-stone. Rocky materials accompanying 
a mineral in this way are called the matrix. The other important 
deposit is in Louisiana. There are many minerals containing 
sulphur but, with the exception of pyrite, these are chiefly impor- 
tant on account of their other constituents. Sulphides of metals, 
such as pyrite FeS 2 , copper pyrites CuFeS 2 , galena PbS, zinc- 
blende ZnS, and sulphates, like gypsum CaS0 4 ,2H 2 0,barite BaS0 4 , 
and celestite SrS0 4 , are fairly plentiful. Sulphur is a constituent 
of the proteins, which are important components of the structure 
of plants and animals. 

Manufacture. — In Sicily, sulphur is obtained by the simple 
process of melting it away from the accompanying volcanic rock 
at a low temperature. The liquid sulphur is allowed to run into 
wooden molds, in which it solidifies in the form of roll sulphur, or 
roll brimstone. To produce the best quality it is subjected to 
distillation from earthenware retorts. When the vapor is led into 
a large brick chamber, it condenses upon the walls and floor at 
first in the form of flowers of sulphur, and later, when the 
chamber becomes heated, as a liquid. 

In Louisiana, the sulphur forms a deposit over half a mile in 
diameter, below 900 feet of clay, quicksand, and rock. It is 
extracted by the Frasch method, by means of borings which 
permit four pipes, one within the other, to reach the deposit. 
Water, previously heated under pressure to 170°, is pumped down 
the two outside pipes (6 and 8 inches in diameter). After time 
has been allowed for the melting of a mass of the sulphur (m.-p. 
114.5°), compressed air is forced down the innermost, one-inch 
pipe. The melted sulphur has twice the specific gravity of the 

264 



SULPHUR AND HYDROGEN SULPHIDE 265 

water in the outer pipes. But the mixture of air and sulphur has 
about the same specific gravity, and so flows freely up the three- 
inch pipe surrounding the air pipe. The element flows into a 
large, wooden enclosure, in which it solidifies, and is practically 
pure sulphur. Each well, until obstructed by collapse of the rock 
and quicksand at the bottom, produces 500 tons a day. 

The greater part of the sulphur of commerce formerly came from 
Sicily, where, in 1898, 447,000 tons were manufactured against 
41,000 tons elsewhere. The whole supply of the United States 
(250,000 tons) is now obtained from Louisiana. The world's 
consumption is over 800,000 tons. 

Physical Properties. — The chief physical peculiarity of 
sulphur is that, instead of appearing in only three familiar physical 
states, like water, it possesses two familiar and perfectly distinct 
solid forms and two different liquid states of aggregation. 

1. Rhombic Sulphur. Native sulphur is yellow, has a sp. gr. 
2.06 and melts at 112.8°. It is almost insoluble in water, but 
dissolves freely in carbon disulphide (41 parts in 100 at 18°). The 
crystals of native sulphur, as w^ell as those obtained by evaporating 
a solution, belong to the rhombic system (Fig. 7, p. 12). Roll 
sulphur and most specimens of flowers of sulphur are the same 
substance although the crystals in their growth have interfered 
with one another, and the mass is crystalline, simply, and not well 
crystallized. This variety is called, from its form, rhombic sul- 
phur. This form is stable below 96°. Above that temperature it 
changes slowly into monoclinic sulphur. 

2. Monoclinic Sulphur. When a large mass of 
melted sulphur solidifies slowly, and the crust is 
pierced and the remaining liquid poured out be- 
fore the whole has become solid, the interior is found 
to be lined with long, transparent needles (Fig. 84) . 
This kind of sulphur is nearly colorless, has a sp. gr. 
1.96, melts at 119.25°, and is in all physical re- 
spects a different individual from rhombic sulphur. 

This variety is named, from the system to which its crystals 
belong, monoclinic sulphur. This form can be kept above 96° 
(transition point, p. 86), but when allowed to cool, it slowly be- 
comes opaque, changing into particles of rhombic sulphur. 




266 COLLEGE CHEMISTKY 

A substance which has two solid states of aggregation and, there- 
fore, two crystalline forms, is said to be dimorphous (two-formed). 

3. S\ and £ M , Vapor. When melted sulphur is heated, it under- 
goes a gradual change, which is especially noticeable near 160°. 
The formerly pale-yellow, mobile liquid (S\) suddenly becomes 
dark-brown in color and so viscous (S M ) that the vessel may be 
inverted without loss of material : Sx <=* S M . The liquid is a mix- 
ture, containing increasing proportions of S M . Beyond 260° the 
viscidity becomes less, and at 444.7° the liquid boils and passes into 
sulphur vapor. 

When ordinary sulphur is raised to the boiling point and then 
allowed slowly to cool, the product is crystalline and soluble in 
carbon disulphide, as before. The change from Sx to S^ is revers- 
ible. But when sulphur is boiled and then suddenly chilled by 
pouring into cold water, it is at first semi-fluid. After several days 
this plastic sulphur, as it is called, becomes hard. It is then found 
to contain rhombic sulphur mixed with 30 per cent of another 
variety of free sulphur, namely &». This part is almost insoluble 
in any solvent. Being without crystalline structure, it is called 
amorphous (Gk., without form) sulphur. Now amorphous bodies 
(see Glass) are always supercooled liquids, that is, liquids still 
existing as such at a temperature at which the solid, crystalline 
form is the stable one. This is simply the S M in a supercooled 
state. When cold, it reverts very slowly to the soluble variety, 
and years are required for the completion of the reversion at 
room temperature. 

Chemical Properties. — At low temperatures and under re- 
duced pressure, the formula of sulphur vapor is Sg. As the tem- 
perature is raised, however, the vapor expands very rapidly, and 
at 800° the molecular weight is 64.2, and the formula therefore S 2 
(p. 117). The formula of dissolved sulphur, as measured by the 
freezing-point method (p. 213), is S 8 . 

Sulphur is an active chemical substance (p. 208). When finely 
divided metals, with the exception of gold and platinum (pp. 60, 
260), are rubbed together with powdered sulphur, union takes 
place and sulphides are produced. Sulphur when heated com- 
bines with great vigor with iron (p. 13), copper, and most of the 
metals. It unites also with many of the non-metals. Thus with 



HYDROGEN SULPHIDE 267 

oxygen it produces sulphur dioxide (p. 31), and even sulphur tri- 
oxide S0 3 . It unites also with chlorine directly. When sulphur 
is treated with oxidizing agents in presence of water, no trace of 
sulphur dioxide (or sulphurous acid) is formed; the only prod- 
uct is sulphuric acid (see p. 289).* 

Uses of Sulphur. — Large quantities of crude sulphur are 
employed for making sulphur dioxide, which is used in the manu- 
facture of sulphuric acid, in bleaching feathers, straw, and wool, 
in preserving dried fruits, and in making alkali sulphites for 
employment in the bleaching industry and in paper-making. The 
manufacture of carbon disulphide also consumes much sulphur. 
Purified sulphur is employed in the manufacture of gunpowder, 
fireworks, matches, and, by combination with rubber, of vulcanite. 
Flowers of sulphur is used in vineyards to destroy fungi, which it 
does by virtue of the traces of sulphuric acid it yields by oxidation. 

Hydrogen Sulphide H 2 S 

This gas is found dissolved in some mineral waters, which in con- 
sequence are known as sulphur waters. It is produced in the de- 
composition of animal matter containing sulphur (proteins), when 
air is excluded. Hence the odor of rotten eggs is due in part to its 
presence. 

Preparation, — 1. Hydrogen and sulphur do not unite percep- 
tibly in the cold. At 310° almost complete union occurs, but about 
168 hours are required for the attainment of equilibrium. 

2. Sulphides of metals, being salts, are acted upon more or less 
easily by dilute acids, and give hydrogen sulphide. Ferrous sul- 
phide, the least expensive of those easily affected, is generally 
used: 

FeS + 2 HC1 £=► H 2 S f + FeCl 2 . 

For hydrochloric acid we may substitute an aqueous solution of 
any active, non-oxidizing acid (see p. 268, last line). A Kipp's 
apparatus (p. 54) is commonly employed. 

* The paragraph on the chemical relations of the element (see end of this 
chapter) should be read at this point. 



268 COLLEGE. CHEMISTRY 

3. Hydrogen sulphide is the invariable product of the extreme 
reduction of any sulphur compound. Thus, it is formed by the 
action of hydrogen iodide upon concentrated sulphuric acid (p. 
201). Even sulphur itself is reduced by dry, gaseous hydrogen 
iodide: 

2HI + S->H 2 S + I 2 . 

Physical Properties. — Hydrogen sulphide is a colorless gas 
with a characteristic odor. When liquefied, it boils at —62°, and 
in solid form melts at —83°. The solubility in water at 10° is 360 
volumes in 100, and becomes less as the temperature is raised. 
The gas can be driven out completely by boiling the solution (cf. 
p. 145). The gas is very poisonous, one part in two hundred of 
air being fatal to mammals. 

Chemical Properties of Hydrogen Sulphide Gas, — When 
heated, the gas dissociates: 

H 2 S <=^ H 2 ~f" S. 

At 310° the decomposition is slight (cf. p. 267), but becomes 
greater at higher temperatures. 

The gas burns in air, forming steam and sulphur dioxide. The 
temperature of the mantle of flame surrounding the gas, as it issues 
from a jet, being far above 310°, the gas in the interior is dissociated 
before it meets with any oxygen. Hence a 
cold dish held across the flame (Fig. 85) re- 
ceives a deposit of free sulphur, and a part of 
the hydrogen also escapes unburnt. It may 
be remarked that dissociation of this kind 
probably precedes the combustion of most 
gaseous compounds (see Flame). 

The metals, down to and including silver 
in the electromotive series, when exposed 
to the gas, quickly receive a coating of sul- 
phide. The tarnishing of silver in the household is probably due 
to a trace of hydrogen sulphide in the illuminating gas which 
escapes from slight leaks in the pipes. That the gas should thus 
behave like free sulphur shows its instability. 

This instability is shown also in the fact that it reduces sub- 




HYDROGEN SULPHIDE 269 

stances, such as sulphur dioxide, which are not affected by free 
hydrogen : 

S0 2 + 2H 2 S->2H 2 + 3S. 

This action takes place in the cold, and much more rapidly when 
the gases are moist than when they are dry (p. 160). Some 
native sulphur is produced by this action, but usually it arises 
from the reduction of gypsum CaS0 4 ,2H 2 to CaS, and libera- 
tion from the sulphide. Sulphur is deposited also when hydrogen 
sulphide undergoes a partial combustion with a restricted supply 
of oxygen, 2H 2 S + 2 — > 2H 2 + 2S, and its formation in nature 
is sometimes to be accounted for in this way. 

A Characteristic of Reduction and Oxidation. — In the 

former of the two actions last mentioned, it will be seen that, while 
the S0 2 was reduced to S, at the same time H 2 S was oxidized (to S). 
In the second action, H 2 S was oxidized to S, and 2 was reduced to 
2H 2 0. It is a characteristic of such actions that one substance is 
oxidized and another reduced: oxidation and reduction always 
occur together, in the same reaction. Here, under hydrogen 
sulphide, we speak of its reducing effect on sulphur dioxide. Under 
sulphur dioxide, however, we should speak of the oxidizing effect 
of the substance on hydrogen sulphide. 

Chemical Properties of the Aqueous Solution of Hydrogen 
Sulphide. — While the gas itself is not an acid, its solution in 
water gives a feeble acid reaction with litmus, and is sometimes 
named hydrosulphuric acid H 2 S, Aq. The conductivity of a iV/10 
aqueous solution is small, and only 0.0007 (0.07 per cent) of the 
substance is ionized: 

H 2 S ±=> H+ + HS" (+± H+ + S=). 

Some S = ions are present. But hydrosulphide-ion HS~, although 
an acid, is less dissociated than is water itself, and the amount of 
sulphide-ion is therefore very small. The salts of hydrosulphide- 
ion, such as NaHS (sodium acid sulphide, see next section), give 
therefore neutral solutions. This behavior is the rule with the 
acid salts of feeble dibasic acids (p. 241). 

As an acid, the solution of hydrogen sulphide may be neutralized 



270 COLLEGE CHEMISTRY 

by bases. For the same reason it enters into double decomposition 
with salts (see next section). 

By the action of oxygen from the air upon an aqueous solution of 
hydrogen sulphide, the sulphur is slowly displaced and appears in 
the form of a fine white powder: 

0» + 2HaS-»2SJ + 2H 2 0. 

This is an action similar to the displacement of ionic bromine by 
free chlorine (p. 259). 

The solution of the gas is a reducing agent, as its action upon 
iodine shows (p. 202). So, also, in presence of an acid, it removes 
oxygen from dichromic acid (produced by the action of an acid 
upon potassium dichromate) : 

K 2 Cr 2 7 + 2HC1 <=> H 2 Cr 2 7 + 2KC1. (1) 

H 2 Cr 2 7 + 6HC1-»4H 2 + 2CrCl 3 (+ 30). , (2) 

(3Q)+3H 2 S->3H 2 Q + 3S. (3) 

Adding: K 2 Cr 2 7 + 8HC1 + 3H 2 S ->2KC1 + 2CrCl 3 + 7H 2 + 3S. 

The first partial equation (c/. p. 194) represents the regular inter- 
action of two ionogens, but the second interaction does not take 
place unless an oxidizable body (here the hydrogen sulphide) is 
present to take possession of the oxygen which it is capable of 
delivering (c/. p. 225). 

The foregoing illustrates a fourth kind of ionic chemical change 
(p. 259), namely that in which a compound ion is formed or decom- 
posed. Here dichromate-ion Cr 2 7 = gives chromic-ion Cr++ + and 
water. For other illustrations see pp. 56, 161, 206, 224, 225, 274. 

Sulphides. — As a dibasic acid (p. 269), hydrogen sulphide 
gives both acid and normal (or "neutral") sulphides, such as 
NaHS and Na 2 S. 

The acid sulphides are obtained by passing the gas in excess into 
solutions of soluble bases: 

H 2 S + NaOH -> H 2 + NaHS, 

and are neutral in reaction. Their negative ion, HS^, is not further 
dissociated (see preceding section). 

By adding to the above-mentioned solution an amount of sodium 
hydroxide equal to that used before, and driving off the water by 



HYDROGEN SULPHIDE 271 

evaporation, the second unit of hydrogen is displaced, and nor- 
mal ("neutral") sodium sulphide is formed: 

NaOH> NaHS t=> Na 2 S + H 2 1 . 

This action is wholly reversed when the dry sodium sulphide is 
dissolved in water, the salt being completely hydrolyzed (p. 197) to 
the acid salt: 

Na2S^2Na+ + S=) _ Hg - 



H 2 0*=>OH- + 



H+j ^ 



The HS~ gives a lower concentration of hydrogen-ion than the 
water, and hence uses up in its formation the ions of hydrogen 
produced by the latter, until an amount of hydroxide-ion equiva- 
lent to half the sodium is formed. The abbreviated equation 
shows this more clearly: 

S= + H+ + OH" -> HS" + OET. 

The solution is therefore strongly alkaline in reaction. In general, 
a normal salt derived from an active base and a weak acid is hydro- 
lyzed to some extent by water and gives an alkaline solution. 

In the abbreviated formulation used above, the union of Na + 
and OH~ to form NaOH is not shown, because it is slight in dilute 
solution and does not affect the result. The union of S = and H+ 
to form HS~ is alone shown, because it is extensive and significant. 
To save space, this plan will be used in future, where the same 
situation exists. 

The soluble acid sulphides are oxidized in aqueous solution by 
atmospheric oxygen: 

2NaSH + 2 -> 2NaOH + 2S. 

The sulphur is not precipitated, but combines with the excess of the 
sulphide, forming polysulphides (see below). Some sodium thio- 
sulphate is produced at the same time. 

The Action of Acids on Insoluble Sulphides. — The inter- 
action of sulphides and acids is itself so important a matter in 
chemistry, and is so similar in theory to many other kinds of 
actions, that special attention should be given to it. The common 
method of preparing hydrogen sulphide from ferrous sulphide 
affords a suitable illustration. 



272 COLLEGE CHEMISTRY 

Since ferrous sulphide is but slightly soluble in water, the action 
proceeds by a rather complex series of equilibria: 

FeS (solid) ±+FeS (dslvd) ±? Fe++ + S= 1 „ Q , A , , w _ a , , 
2HC1 fc?2Cl"+2H+ J ^ 2 (dslvd)^±H 2 S (gas). 

It will be seen that a number of reversible changes are involved, 
and the question is, why does the reaction proceed forward, as it 
does? To answer this question, a consideration of each of the 
equilibria, separately, is required. 

1. The dissolved hydrogen sulphide is very feebly ionized, and 
maintains a smaller concentration of sulphide-ion S- than does 
ferrous sulphide, in spite of the comparative insolubility of the 
latter. Hence, the S— formed from the FeS is continuously re- 
moved by union with the hydrogen-ion furnished by the acid, 
S = + 2H + ±+ H 2 S, and all the other equilibria are constantly dis- 
placed forward on this account. The action is therefore, in 
essence, like neutralization (p. 254). 

2. The union of S~ and 2H+ depends on the magnitude of the 
product of their concentrations (p. 184), [S=] X [H+] X [H+], or 
[S = ] X [H+] 2 . Hence, although [S— ] is minute, on account of the 
insolubility of FeS, [H + ] is large on account of the great dissocia- 
tion of the HC1 and the fact that a strong solution of the acid can 
be used. Thus the product may be large enough for the purpose. 

3. When a still more insoluble sulphide, like cupric sulphide 
CuS is employed, the concentration of the sulphide-ion [S = ] is too 
small to play its part and the action makes almost no progress. In 
this case, a concentration of H+, sufficient to raise the product to 
the necessary value, cannot be obtained with any acid. 

4. The fact that hydrogen sulphide is fairly soluble (3.6 vols. : 1 
vol.) hinders the action. It prevents that free escape of one prod- 
uct which is so constantly a factor in promoting reversible chemical 
changes. Thus, if cadmium sulphide CdS, which lies between 
ferrous and cupric sulphides, in solubility, is employed along with 
rather dilute hydrochloric acid, a concentration of hydrogen sul- 
phide sufficient to stop the action accumulates before the liquid is 
saturated with the gas, and the latter can begin to escape. There 
are then two ways of making this action continuous. Either 
stronger hydrochloric acid, giving a higher concentration of H+ 
may be used to force the formation of more H 2 S (by union of 2H + 



HYDROGEN SULPHIDE 273 

and S=), or the reverse action, due to accumulation of H 2 S (dslvd), 
may be diminished mechanically by leading air through the mix- 
ture (p. 129) and so removing the hydrogen sulphide as fast as it 
is formed. Either plan will cause complete interaction with the 
cadmium sulphide. 

Classification of Insoluble Sulphides. — In analytical 
chemistry, advantage is taken of the different solubilities of the 
sulphides, for the purpose of identifying the metallic elements, and 
of separating mixtures containing several such elements. Three 
classes are distinguished. 

1. The sulphides of silver, copper, mercury, and some other 
metals are exceedingly insoluble, and, therefore, do not interact 
with dilute acids as does ferrous sulphide (p. 271). These may 
therefore be made by leading hydrogen sulphide into solutions of 
their salts: 

CuS0 4 + H 2 S fc? CuS | + H 2 S0 4 . 

The acid produced has scarcely any effect upon the sulphide, and 
almost no reverse action is observed. In this action the sulphide- 
ion is the active substance and, by its removal, all the equilibria 
are displaced forwards. 

2. The sulphides of iron, zinc, and certain other metals are insol- 
uble in water, but not so much so as the last class. Hence they are 
decomposed by dilute acids, and the reverse of the above action 
takes place almost completely. These sulphides must therefore be 
made, either by combination of the elements, or by adding a soluble 
sulphide to a solution of a salt : 

FeS0 4 + (NH4) 2 S ^ FeS J + (NH^O* 

No acid is produced in this sort of interaction, and the considerable 
insolubility of the sulphide of iron or zinc in water renders the 
change nearly complete. 

3. The sulphides of barium, calcium, and some other metals 
(q.v.), although insoluble in water, are hydrolyzed by it, and give 
soluble products, the hydroxide and hydrosulphide : 

2CaS + 2H 2 fc* Ca(OH) 2 + Ca(SH) 2 . 

They may be prepared by direct union of the elements, and from 
the sulphates by reduction with carbon. But they are not pre- 
cipitated by hydrogen sulphide or ammonium sulphide. 



274 COLLEGE CHEMISTRY 

Poly sulphides. — When sulphur is shaken with a solution of a 
soluble sulphide or acid sulphide, such as sodium sulphide, it dis- 
solves, and evaporation of the solution leaves residues, varying in 
composition from Na 2 S 2 to Na 2 S 5 . These appear to be mixtures 
composed mainly of Na 2 S and Na 2 S 4 . 

When an acid is poured into sodium polysulphide solution, 
minute spherules of rhombic sulphur are precipitated: 

NaA + 2HC1 -> 2NaCl + H 2 S | + 3S | . 

The Chemical Relations of the Element Sulphur, — In 

combination with metals and hydrogen, sulphur is bivalent, form- 
ing compounds like H 2 S, FeS, CuS, and HgS. In combination with 
non-metals, however, the valence is frequently greater, the maxi- 
mum being seen in sulphur trioxide, where the sulphur is sexivalent. 
Its oxides are acid-forming, and it is, therefore, a non-metal. 

Exercises. — 1. How could the decomposition of hydrogen sul- 
phide at 310° be rendered, (a) more complete, (6) less complete? 
Would the percentage decomposed be affected, (a) by reducing the 
pressure, (6) by mixing the gas with an indifferent gas? 

2. What are the relative volumes of the gases (p, 150) in the 
action of, (a) hydrogen iodide and sulphur, (b) hydrogen sulphide 
and sulphur dioxide? 

3. To what classes of ionic actions (p. 259) do the interactions of 
hydrogen sulphide solution with, (a) oxygen (p. 270), (b) sodium 
hydroxide (p. 270), (c) iodine (p. 202) belong? 

4. Show which actions on the pages referred to on p. 270 illus- 
trate the fourth kind of ionic chemical change, and how they do so? 

5. Why is normal sodium sulphide only half hydrolyzed by 
water? 

6. Formulate completely, after the model on p. 252, the actions 
of (a) hydrogen sulphide and cupric sulphate solution; (6) am- 
monium sulphide and ferrous sulphate. In each case explain 
which equilibrium determines the direction of the action. 



CHAPTER XXI 

THE OXIDES AND OXYGEN ACIDS OF SULPHUR 

The only important oxides of sulphur are the dioxide S0 2 and 
the trioxide S0 3 . They are the anhydrides (p. 94) of sulphurous 
acid H 2 S0 3 and of sulphuric acid H 2 S0 4 , respectively. 

The Preparation of Sulphur Dioxide S0 2 * — 1. When sulphur 
burns in air or oxygen, sulphur dioxide is produced (p. 32). 2. 
The larger part of the sulphur dioxide used in commerce is probably 
obtained by the roasting (calcining) of sulphur ores. Pyrite FeS 2 , 
for example, on account of the large amount of sulphur which it 
contains, can be burnt in a suitable furnace : 

4FeS 2 + 110 2 -> 2Fe 2 3 + 8S0 2 1 . 

The gas, although mixed with great amount of nitrogen which 
entered as part of the air, can be used to make sulphuric acid. 

It should be noted, in passing, that heating and roasting or cal- 
cining are distinct processes in chemistry. Roasting or calcining 
always assumes the access of the air and employment of its oxygen; 
heating, in the absence of modifying words, assumes the exclusion 
or the chemical indifference of the air. 

3. In the laboratory, a steady stream of the gas is obtained by 
allowing hydrochloric acid to drop upon solid sodium acid sul- 
phite, or concentrated sulphuric acid to trickle into a 40 per cent 
solution of the same salt (Fig. 24, p. 54) : 

HC1 + NaHS0 3 ±+ NaCl + H 2 S0 3 *=* H 2 + SOrf. 

The sulphurous acid, being very unstable, decomposes spontane- 
ously into water and sulphur dioxide, and the latter escapes when 
sufficient water for its solution is not present. 

4. Sulphur dioxide can also be made by the reduction of con- 
centrated sulphuric acid by copper at a high temperature. A part 

275 



276 COLLEGE CHEMISTRY 

of the acid loses oxygen to form water with the hydrogen of 
another molecule: 

Partial 1 : H 2 S0 4 -> H 2 + S0 2 (+ 0). 

Partial 2: (O) + H 2 S0 4 + Cu -> H 2 + CuS0 4 . 

2H 2 S0 4 + Cu -> 2H 2 + S0 2 + CuS0 4 . 

Some easily oxidized non-metals, such as carbon and sulphur, act 
in the same way, C + 2H 2 S0 4 -» 2H 2 + 2S0 2 + C0 2 . 

Making Equations by Positive and Negative Valences. — 

Equations like the foregoing can be constructed also by assuming 
that each element in a compound is either positive or negative, 
and by marking the valences accordingly (for details, see p. 322). 
Thus, in sulphuric acid, we have 2H+ (positive, univalent) and 40= 
(each bivalent and negative). Since the numbers of positive and 
negative valences must be equal, and we have 2©* and 80, it 
follows that the sulphur carries 6 © , Sttt. 

Now when, in making the experiment, we find the products S0 2 
and CuS0 4 , we may infer that the hydrogen formed water. We 
infer, also, that to obtain two compounds containing sulphur, at 
least 2H 2 S0 4 was required. We then note that the S in S0 2 is 
quadrivalent. Hence Sttt became Stt and 2© were released. 
The metallic copper used was free and without valence, and be- 
came CuS0 4 , in which it is CU++. It obtained the 2© from the 
sulphur. The action can therefore be analyzed as follows: 

[2H+ + Sttt + 40=] r* [Stt + 20=]+ [2H+ + 0=] + [0= + 2©] 
First H 2 S0 4 S0 2 H 2 Balance 

The second H 2 S0 4 gives [2H+ + S0 4 =]. The Cu takes the 2© 
giving Cu ++ , and this with the S0 4 = gives CuS0 4 . The 2H+ takes 
the 0= from the balance, giving H 2 0. Thus, the whole balance is 
used and the products are accounted for. The equation must 
therefore be: 

2H 2 S0 4 + Cu -> S0 2 + 2H 2 + CuS0 4 . 

It will be noted that the two molecules of sulphuric acid play 
different roles. Only one of them is used in oxidizing. 

* The signs © and stand for quantities of electricity equal to those 
carried by one equivalent of an ionic substance, and therefore required for 
its discharge and liberation. 



THE OXIDES AND OXYGEN ACIDS OF SULPHUR 



277 



Similarly, with sulphuric acid and carbon, the same analyzed 
equation applies. The carbon gives C0 2 . Thus, the carbon goes 
from C° to Ctt. To obtain the 4©, 2H 2 S0 4 is required (equation 
above). Hence, 

2H 2 S0 4 + C -> C0 2 + 2S0 2 + 2H 2 0. 

When hydrogen sulphide is led through concentrated sulphuric 
acid, the latter is reduced to sulphur dioxide, and the former is 
oxidized, giving free sulphur (p. 270) : 

2H+ + S= + 2© -> 2H+ + S°|. 

Since this action requires 2©, and sulphuric acid in giving S0 2 
delivers 2©, it follows that 1H 2 S0 4 will decompose 1H 2 S: 
H 2 S0 4 + H 2 S -> 2H 2 + S°| + S0 2 . 

Finally, when HI with sulphuric acid (p. 201) gives free iodine 
(1°), and H 2 S (2H + + S=), evidently Sffi in sulphuric acid gives 
up 8©, becoming S~: 

[2H+ + Sttt + 40=] -> 2H+ + S= + 40= + 8© 
and [H+ + I"] + -» H+ + 1°. 

Evidently, 1H 2 S0 4 giving 8© will interact with 8HI, changing 
8I~ into 81°. Hence, 

H 2 S0 4 + 8HI -> 4H 2 + H 2 S + 81°. 

The reader should practice the use of this method by making the 
equations for the actions of zinc (p. 268 giving 
hydrogen sulphide) and of hydrogen bromide 
(p. 196) upon sulphuric acid. 

Physical and Chemical Properties. — 

Sulphur dioxide is a gas possessing a pene- 
trating and characteristic odor. This is fre- 
quently spoken of as the "odor of sulphur," 
but it should be remembered that sulphur 
itself has scarcely any smell at all. The 
weight of the G.M.V. of the gas (65.54 g.) 
shows it to be more than twice as heavy as air. 
By means of a freezing mixture of ice and salt (Fig. 86), the gas 
is easily condensed in a (J -tube to a transparent mobile fluid, which 




Fig. 86. 



278 COLLEGE CHEMISTRY 

boils at —8°. At 20°, the liquid gives a vapor pressure of only- 
Si atmospheres, so that the liquid is handled and sold in glass 
syphons or in sealed tin cans. The solubility of the [gas in water 
is 5000 volumes in 100. The liquid is completely freed from the 
gas by boiling (cf. p. 145). 

As regards chemical properties, sulphur dioxide is stable (p. 93). 

It unites with water to form sulphurous acid H2SO3, which is 
unstable, and exists only in solution. 

Since the maximum valence of sulphur is 6, sulphur dioxide, in 
which but four of the valences of sulphur are used, is unsaturated. 
It is therefore still able to combine directly with suitable elements, 
such as chlorine and oxygen. When it is mixed with chlorine in 
sunlight, a liquid, sulphuryl chloride SO2CI2 is produced. 

Liquefied sulphur dioxide is employed for bleaching straw, wool, 
and silk (see p. 289). As a disinfectant it has been displaced to a 
large extent by formaldehyde. 

The Liquefiability of Gases. — It will assist us in recalling 
which gases are hard to liquefy and which easy, if we memorize 
the fact that Faraday (from 1823 to 1845) liquefied most of the 
familiar gases and failed only with three, namely hydrogen (c.t. 
-242°), oxygen (c.t. -113°), and nitrogen (c.t. -146°). These, 
with nitric oxide NO (c.t. —93.5°), carbon monoxide CO (c.t.— 40°), 
methane CH4 (c.t. —99°), and the six inert gases (pp. 335-337), 
are the ones which have low critical temperatures (cf. p. 78) and 
are difficult to liquefy. 

Of the gases we have studied, the ones which are more or less 
easily liquefied are: hydrogen chloride (c.t. +52°), bromide, and 
iodide, chlorine (c.t. +141°), ozone, hydrogen sulphide (c.t. + 100°), 
sulphur dioxide (c.t. +154°). 

The Solubilities of Gases. — For the purpose of remember- 
ing the solubilities of gases in water, it is convenient to divide the 
gases into three classes. The following are the ones we have 
studied: 

1. Slightly soluble: Oxygen (4 vol. : 100 at 0°), hydrogen 
(2 : 100 at 0°). 

2. Soluble: Chlorine (260 vol. : 100 at 10°), hydrogen sul- 
phide (440 : 100 at 0°). 



THE OXIDES AND OXYGEN ACIDS OF SULPHUR 279 

3. Very soluble: Hydrogen chloride (505 vol. : 1 at 0°), bro- 
mide (404 : 1) and iodide (1570 : 1), sulphur dioxide (69 : 1 at 0°). 

Preparation of Sulphur Trioxide SO3, — Although the for- 
mation of sulphur trioxide is accompanied by the liberation of much 
heat, sulphur dioxide and oxygen, whether cold or warm, unite 
very slowly. Ozone, however, combines with the former readily. 

The interaction of sulphur dioxide and oxygen is hastened by 
finely divided platinum, which remains itself unchanged and simply 
acts as a catalytic agent. The contact 'process, as this is called, 
has been rendered available for the commercial manufacture of 
sulphur trioxide by Knietsch (1901). At 400°, the temperature 
used, 98-99 per cent of the materials unite. 

2 + 2S0 2 -> 2S0 3 + 2 X 22,600 cal. 

Below 400°, the union is too slow. Above 400°, the reverse action 
is strengthened (Van't Hoff's law, p. 188), and the union is too 
incomplete. The vaporous product is condensed by being led 
into 97-99 per cent sulphuric acid, and the concentration of the 
liquid is constantly maintained at this point by the regulated in- 
flux of water. The sulphur dioxide is obtained by calcining ores 
(p. 275). These contain impurities which must be removed very 
thoroughly. Dust from the roasting and oxide of arsenic, which 
are present, will otherwise " poison" the contact agent (platinum 
or ferric oxide) and soon almost stop the union. 

The process may be illustrated by placing some platinized 
asbestos* in a tube (Fig. 66, p. 156), which is gently warmed, and 
introducing oxygen and sulphur dioxide through the limbs of the 
Y-tube. Dense fumes appear at the exit (see next section). 

Formerly sulphur trioxide was obtained by the distillation of 
impure ferric sulphate, Fe 2 (S0 4 )3 — » Fe20 3 + 3S0 3 . 

Physical and Chemical Properties. — Sulphur trioxide is a 
volatile liquid (b.-p. 46°). The crystals, obtained by cooling, melt 
at 14.8°. It fumes strongly when exposed to the air, in conse- 
quence of the union of the vapor with moisture and the production 
of minute drops of sulphuric acid. A white crystalline variety, 

* Asbestos, dipped in a solution of chloroplatinic acid and heated in the 
Bunsen flame: H 2 PtCle -» Pt + 2HC1 1 + 2C1 2 T . 



280 COLLEGE CHEMISTRY 

closely resembling asbestos in appearance, is the more familiar 
form of the substance, which is dimorphous (p. 266) . 

As to chemical properties, the vapor of sulphur trioxide dissociates 
into sulphur dioxide and oxygen (400°, 2%; 700°, 40%). 

Sulphur trioxide is not itself an acid, but it is the anhydride 
of sulphuric acid. When placed in water it unites vigorously, 
causing a hissing noise due to the steam produced by the heat of 
the union. 

Just as sulphur trioxide unites with water to give hydrogen 
sulphate, so it combines vigorously with many oxides of metals, 
producing the corresponding sulphates: 

H 2 + S0 3 *=► H 2 S0 4 , CaO + S0 3 -> CaS0 4 . 

The union of an oxide of a non-metal with the oxide of a metal, in 
this fashion, is a general method of obtaining salts (c/. p. 261). 

Oxygen Acids of Sulphur. — Sulphurous and sulphuric acids 
have been mentioned frequently already. Next to them in im- 
portance come thiosulphuric acid and persulphuric acid. The 
compositions of the acids show their relationships: 

Hyposulphurc-ws acid, H2S2O4. Sodium hyposulphite, Na 2 S 2 4 . 



Sulphurous acid, 


H 2 S0 3 . 


Sodium sulphite, 


Na 2 S0 3 . 


Sulphuric acid, 


H 2 S0 4 . 


Sodium sulphate, 


Na 2 S0 4 . 


Thiosulphuric acid, 


H 2 S 2 03. 


Sodium thiosulphate, 


Na 2 S 2 3 . 


Persulphuric acid, 


H 2 S 2 0s. 


Sodium persulphate, 


Na 2 S 2 8 . 



Thiosulphuric acid (Gk. Oeiov, sulphur) is so named because it 
contains one unit of sulphur in place of one of the units of oxygen 
of sulphuric acid. Note that when the names of the acids end in 
ous and ic, the names of the salts end in ite and ate, respectively. 
Besides the above we have also the polythionic acids, namely: 
dithionic acid H 2 S 2 06, trithionic acid H 2 Ss06, tetrathionic acid 
H3S4O6, and pentathionic acid E^SsOe. 

Sulphuric Acid H 2 S0 4 

Although salts of sulphuric acid, such as calcium sulphate CaSC>4, 
are exceedingly plentiful in nature, the preparation of the acid by 
chemical action upon the salts is not practicable. The sulphates, 
indeed, interact with all acids, but the actions are reversible. The 
completion of the action by the plan used in making hydrogen 



THE OXIDES AND OXYGEN ACIDS OF SULPHUR 281 

chloride (p. 142), involving the removal of the sulphuric acid by 
distillation, would be difficult on account of the involatility of this 
acid. It boils at 330°; and suitable acids, less volatile still, which 
might be used to liberate it, do not exist. We are therefore com- 
pelled to build up sulphuric acid from its elements. 

The union of sulphur dioxide and oxygen by the contact process, 
and combination of the trioxide with water (p. 279), is the best 
method for making a highly concentrated acid. For obtaining 
ordinary "oil of vitriol," however, the "chamber process" is still 
used extensively. 

Chemistry of the Chamber Process, — The gases, the inter- 
actions of which result in the formation of sulphuric acid, are: 
water vapor, sulphur dioxide, nitrous anhydride N2O3* (see index), 
and oxygen. These are obtained, the first by injection of steam, 
the second usually by the burning of pyrite, the third from nitric 
acid HNO3, and the fourth by the introduction of air. The gases 
are thoroughly mixed in large leaden chambers, and the sulphuric 
acid forms droplets which fall to the floors. In spite of elaborate 
investigations, instigated by the extensive scale upon which the 
manufacture is carried on and the immense financial interests 
involved, some uncertainty still exists in regard to the precise 
nature of the chemical changes which take place. According to 
Lunge, supporting the view first suggested by Berzelius, the greater 
part of the product is formed by two successive actions, the first 
of which yields a complex compound that is decomposed by excess 
of water in the second: 

O-H 

h 2 o + 2so 2 + N2O3 + o 2 -> 2so 2 : (i) 

x O-NO 

The group —NO, nitrosyl, is found in many compounds. Here, if 
it were displaced by hydrogen, sulphuric acid would result. Hence 
this compound is called nitrosylsulphuric acid: 

/O-H OH 

2SO2C +H 2 0^±2S0 2 . + N 2 3 . (2) 

X 0-NO x OH 

* This gas is unstable, breaking up in part into nitric oxide NO and nitro- 
gen tetroxide N0 2 : N 2 3 <=* NO + N0 2 . In this process, however, the mix- 
ture behaves as if it were all N 2 3 , and so only nitrous anhydride is named in 
this connection. 



282 COLLEGE CHEMISTRY 

The equations (1) and (2) are not partial equations for one inter- 
action, but represent distinct actions which can be carried out 
separately. In a properly operating plant, indeed, the nitrosyl- 
sulphuric acid is not observed. But when the supply of water is 
deficient, white "chamber crystals," consisting of this substance, 
collect on the walls. 

The explanation of the success of this seemingly roundabout 
method of getting sulphuric acid is as follows: The direct union of 
sulphur dioxide and water to form sulphurous acid is rapid, but the 
action of free oxygen upon the latter, 2H 2 S0 3 + 2 — > 2H 2 S0 4 , is 
exceedingly slow. Reaching sulphuric acid by the use of these two 
changes, although they constitute a direct route to the result, is not 
feasible in practice. On the other hand, both of the above actions, 
(1) and (2), happen to be much more speedy, and so, by their use, 
more rapid production of the desired substance is secured at the 
expense of a slight complexity. 

The progress of the first action is marked by the disappearance 
of the brown nitrous anhydride and, on the introduction of water, 
the completion of the second stage results in the reproduction of 
the same substance. The nitrous anhydride takes part a large 
number of times in these changes, and so facilitates the conversion 
of a great amount of sulphur dioxide, oxygen, and water into sul- 
phuric acid, without much diminution of its quantity. Some is 
lost, however. 

The loss of nitrous anhydride is made good by the introduction 
of nitric acid vapor into the chamber. This acid is made from con- 
centrated sulphuric acid and commercial sodium nitrate NaN0 3 : 

|NaN0 3 + H 2 S0 4 <=± HN0 3 T + NaHS0 4 . 

On account of the volatility of the nitric acid, a moderate heat is 
sufficient to remove it from admixture with the other substances, 
and its vapor is swept along with the other gases into the apparatus. 
The initial action which the nitric acid undergoes: 

H 2 + 2S0 2 + 2HN0 3 -> 2H 2 S0 4 + N 2 3 , 

may be written, to show the anhydride of nitric acid : 

H 2 + 2S0 2 + H 2 0,N 2 5 -> 2H 2 S0 4 + N 2 3 . 

The two molecules of water, one actually, the other potentially, 
present, with the two molecules of sulphur dioxide, can furnish two 



THE OXIDES AND OXYGEN ACIDS OF SULPHUR 



283 



molecules of sulphurous acid (H 2 S0 3 ). The N 2 5 in passing to the 
condition N 2 3 gives up the two units of oxygen required to con- 
vert this sulphurous acid into sulphuric acid. 

Details of the Chamber Process. — The sulphur dioxide is 
produced in a row of furnaces A (Fig. 87). When good pyrite is 
used, the ore burns unassisted (p. 275), while impure pyrite and 
zinc-blende ZnS have to be heated artificially to maintain the com- 
bustion. The gases from the various furnaces pass into one long 




Fig. 87. 

dust-flue, in which they are mingled with the proper proportion of 
air, and deposit oxides of iron and of arsenic, and other materials 
which they transport mechanically. From this flue they enter the 
Glover tower G, in which they acquire the oxides of nitrogen. 
Having secured all the necessary constituents, excepting water, the 
gases next enter the first of the lead chambers, large structures 
lined completely with sheet lead. These measure as much as 
100 X 40 X 40 feet, and have a total capacity of 150,000 to 200,000 
cubic feet. As the gases drift through these chambers they are 
thoroughly mixed, and an amount of water considerably in excess 
of that actually required is injected in the form of steam at various 
points. The acid, along with the excess of water, condenses and 



284 COLLEGE CHEMISTRY 

collects upon the floor of the chamber, while the unused gases, 
chiefly nitrous anhydride and nitrogen, the latter derived from the 
air originally admitted, find an exit into the Gay-Lussac tower L. 

This is a tower about fifty feet in height, filled with tiles, over 
which concentrated sulphuric acid continually trickles. The object 
of this tower, to catch the nitrous anhydride and enable it to be 
reemployed in the process, is accomplished by a reversal of action 
(2) above. The acid which accumulates in the vessel at the bottom 
of this tower contains the nitrosylsulphuric acid, and by means of 
compressed air is forced through a pipe up to a vessel at the top of 
the Glover tower G. When this " nitrous vitriol" is mixed with 
dilute sulphuric acid from a neighboring vessel, by allowing both to 
flow down into the tower, the nitrous anhydride is once more set 
free by the interaction of the water in the dilute acid (action (2)). 
The Glover tower is filled with broken flint or tiles, and the heated 
gases from the furnace acquire in it their supply of nitrous anhy- 
dride. Their high temperature causes a considerable concentra- 
tion of the diluted sulphuric acid as it trickles downward. The 
acid, after traversing this tower, is sufficiently strong to be used 
once more for the absorption of nitrous anhydride. 

To replace the part of the nitrous anhydride which is inevitably 
lost, fresh nitric acid is furnished by small open vessels N, contain- 
ing sodium nitrate and sulphuric acid, placed in the flues of the 
pyrite-burners. About 4 kg. of the nitrate are consumed for every 
100 kg. of sulphur. 

The acid which accumulates upon the floors contains but 60 to 
70 per cent of sulphuric acid, and has a specific gravity of 1.5-1.62. 
The excess of water is needed to facilitate the second action. It is 
required also in order that the acid upon the floor may not after- 
wards absorb and retain the nitrous anhydride, for this substance 
combines with an acid containing more than 70 per cent of hydro- 
gen sulphate. 

This crude sulphuric acid is applicable directly in some chemical 
manufactures, such as the preparation of superphosphates (q.v.). 
Concentration is effected by evaporation in pans lined with lead, 
which are frequently placed over the pyrite-burners in order to 
economize fuel. The evaporation in lead is carried on until a 
specific gravity 1.7, corresponding to 77 per cent concentration, is 
reached. Up to this point the sulphate of lead formed by the 



THE OXIDES AND OXYGEN ACIDS OF SULPHUR 285 

action of the sulphuric acid produces a crust which protects the 
metal from further action. When a stronger acid is required, 
the water is driven out by heating the sulphuric acid in vessels of 
glass or platinum, or even of cast iron. Iron acts upon dilute 
sulphuric acid, displacing the hydrogen-ion, but not upon concen- 
trated sulphuric acid, which is not ionized. Commercial sulphuric 
acid, oil of vitriol, has a specific gravity 1.83-1.84, and contains 
about 93.5 per cent of hydrogen sulphate. 

Physical Properties. — Pure hydrogen sulphate has a sp. gr. 
1.85 at 15°. When cooled, it crystallizes (m.-p. 10.5°). At 150°- 
180° the acid begins to fume, giving off sulphur trioxide. It 
boils at 330°, but loses more sulphur trioxide than water and finally 
yields an acid of constant (p. 145) boiling-point (338°) and con- 
stant composition (98.33 per cent). The heat of solution (p. 125) 
of hydrogen sulphate is very great (39,170 cal.). The solution is 
thus much more stable (i.e., it contains much less energy) than the 
pure substance, and hence the latter absorbs water greedily. 

Commercial sulphuric acid is impure. It contains, for example, 
lead sulphate, which appears as a precipitate when the acid is 
diluted, as well as arsenic trioxide and oxides of nitrogen in com- 
bination. 

Chemical Properties and Uses of Hydrogen Sulphate. — 

1. The compound is not exceedingly stable, for dissociation into 
water and sulphur trioxide begins far below the boiling-point. 
The vapor of the acid boiling at 338° contains 30 per cent of 
H 2 + S0 3 , which recombine when the vapor is condensed. The 
dissociation is practically complete at 416°, as is shown by the 
density of the vapor. When raised suddenly to a red heat it is 
broken up completely into water, sulphur dioxide, and oxygen. 

2. When sulphur trioxide is dissolved in hydrogen sulphate, di- 
sulphuric acid H2S2O7, a solid compound, is obtained. Hydrogen 
sulphate containing 80 per cent of disulphuric acid is known as 
"oleum," and is employed in chemical industries. The salts of 
disulphuric acid may be made by strongly heating the acid sul- 
phates, for example: 

2NaHS0 4 <=± NaAOy + H 2 OT. 



286 COLLEGE CHEMISTRY 

In view of this mode of preparation by the aid of heat, they are 
frequently known as pyrosulphates (Gk. nvp, fire). When they are 
dissolved in water, the acid sulphates are reproduced. 

3. With salts which it does not oxidize (see below), hydrogen sul- 
phate reacts by double decomposition and sets free the correspond- 
ing acid. Where the new acid is volatile, as in the case of hydrogen 
chloride (p. 142), we are furnished with one of the cheapest means 
of preparing acids. Since hydrogen sulphate is dibasic (p. 245), 
it forms both acid and normal salts, such as NaHS0 4 and Na 2 S0 4 . 
The acid sulphates are called also bisulphates, because they con- 
tain twice as large a proportion of S0 4 to Na, and require twice as 
much sulphuric acid for their preparation as do the neutral sul- 
phates. 

4. Sulphuric acid combines vigorously with "water to form at 
least one rather stable hydrate, H 2 S0 4 ,H 2 (m.-p. 8°). On this 
account, sulphuric acid is able to take the elements of water from 
compounds containing hydrogen and oxygen, especially those con- 
taining these elements in the proportion 2H : 0. Thus paper, 
which is largely cellulose (CeHioOs)*, wood which contains much 
cellulose, and sugar C12H22O11 are charred by it, and carbon is 
set free: 

C^H^Oi! -> 12C + 11H 2 0. 

For the same reason, sulphuric acid is used in drying gases with 
which it does not interact. 

5. On account of the large quantity of oxygen which hydrogen 
sulphate contains, and its instability when heated, it behaves as 
an oxidizing agent. This property has already been illustrated in 
connection with the action of the acid upon carbon, sulphur, and 
copper (p. 276), hydrogen iodide (p. 201), and hydrogen bromide 
(p. 196). The sulphuric acid is in consequence reduced to sulphur 
dioxide, and even to free sulphur or hydrogen sulphide. The 
metals, from the most active down to silver (p. 260), are capable 
of reducing it, the sulphates* being formed. The more active 
metals, like zinc, reduce it to hydrogen sulphide (p. 277), the less 

* Note that the sulphates, and not the oxides of the metals are produced. 
Oxides of metals could not be formed in concentrated sulphuric acid, because 
they interact with the latter much more vigorously than do the metals, to 
give the sulphates (cf. p. 146). 






THE OXIDES AND OXYGEN ACIDS OF SULPHUR 287 

active, like copper, give sulphur dioxide (p. 276). Hydrogen is 
not liberated, because no hydrogen-ion is present in concentrated 
sulphuric acid. Gold and platinum alone do not interact with it. 
Free hydrogen itself is oxidized to water when passed into hydrogen 
sulphate at 160°: S0 2 (OH) 2 + H 2 -> S0 2 + 2H 2 0. 

Concentrated sulphuric acid is used in almost all chemical in- 
dustries: for example, to give sodium sulphate, as a stage in the 
Le Blanc process for the manufacture of soda; in the refining of 
petroleum; in the manufacture of fertilizers, such as superphos- 
phate; in the preparation of nitroglycerine and gun-cotton, where 
it assists the action by removing water; and in the production of 
coal-tar dyes. 

Chemical Properties of Aqueous Hydrogen Sulphate, — 

The solution of sulphuric acid H 2 S0 4 ,Aq is a mixture, whose com- 
ponents are: undissociated molecules H 2 S0 4 , hydrogen-ion H + , 
hydrosulphate-ion HS0 4 ~, and sulphate-ion S0 4 =\ The chemical 
properties shown by the solution are those of one or other of these 
components, according to circumstances. 

Except in concentrated solutions (normal or stronger) the oxidiz- 
ing effects of the undissociated, molecular substance are not 
encountered. 

The presence of hydrogen-ion is shown by all its usual properties 
(p. 246). 

Sulphate-ion S0 4 — , which is found also in solutions of all neutral 
and acid sulphates, unites with all positive ions. The product, 
when insoluble, appears as a precipitate. The introduction of 
barium ions, for example, by adding a solution of barium nitrate 
or chloride, is employed as a test: 

Ba++ + S0 4 =^BaS0 4 |. 

Since there are other barium salts which are insoluble in water (see 
Table of Solubilities), but no common ones which are not decom- 
posed by acids, dilute nitric acid is first added to the solution 
supposed to contain the sulphate-ion. The other ions, even if 
present, then give no precipitate with barium-ion. 

Dilute sulphuric acid is used for many purposes. Thus, it 
forms the liquid in the lead storage battery, and is employed for 
cleaning sheet iron before tinning and galvanizing. 



288 COLLEGE CHEMISTRY 

Sulphates, — The acid sulphates, known also as bisulphates, 

(see p. 286), may be produced either by adding to dilute sulphuric 
acid half an equivalent of a base, and evaporating: NaOH + 
H 2 S0 4 £5 H 2 + NaHS0 4 , or by actions in which another acid is 
displaced, as in making hydrogen chloride (p. 141). These salts 
are acid in reaction, as well as in name (c/. p. 269), because HS0 4 ~, 
although a weak, is not a feeble acid. When heated, they yield 
pyrosulphates (p. 286). 

The normal (or neutral) sulphates are obtained by complete 
neutralization and evaporation, or by the second of the above 
methods when a sufficient amount of the salt and a higher tempera- 
ture are used: 

NaCl + NaHS0 4 *± Na 2 S0 4 + HClt- 

They may also be made by precipitation, by oxidation of a sulphide 
at a high temperature, PbS + 20 2 — ■> PbS0 4 , or by addition of 
sulphur trioxide to the oxide of a metal (p. 280). 

Normal sulphates of the heavy metals decompose at a red heat, 
some giving off sulphur trioxide (p. 279), others sulphur dioxide 
and oxygen. The sulphates of the more active metals and of lead, 
however, are not affected by heating. 



Other Acids of Sulphur 

Sulphurous Acid H 2 S0 3 , Aq. — This term is applied to the 
solution of sulphur dioxide in water. A portion of the sulphur 
dioxide remains dissolved physically, while another portion is in 
combination with the water, forming sulphurous acid. This in 
turn is ionized, and chiefly ? after the manner of the weaker dibasic 
acids, into two ions, H + and HS0 3 ~~. A little S0 3 = is formed from 
the latter. 

Properties of Sulphurous Acid. — The acid is so unstable that 
it cannot be obtained excepting in solution in water. Chemically it 
is a comparatively weak acid. As a reducing agent, it is slowly 
oxidized to sulphuric acid by free oxygen. Sugar and glycerine 
act as negative contact agents and make the oxidation much slower. 
It is oxidized more rapidly by oxidizing agents. Thus, when free 



THE OXIDES AND OXYGEN ACIDS OF SULPHUR 289 

halogens are added to the solution (cf. p. 161), sulphuric acid and 
the hydrogen halide are formed: 

H 2 S0 3 + HIO *± H 2 S0 4 + HI. 

Hydrogen peroxide, potassium permanganate, and other oxidizing 
agents convert the substance into sulphuric acid likewise. 

Sulphurous acid has the power of uniting directly with many 
organic coloring matters and, since the products of this union are 
usually colorless, it is employed as a bleaching agent. It is 
especially useful with chemically reactive materials like silk, wool, 
and fragile structures like straw, which are likely to be destroyed 
if bleaching powder is used. The compounds thus formed are 
unstable, and lose the sulphurous acid again. Hence, wool yellows 
with age, and straw hats lose their whiteness. As a disinfectant 
it acts, perhaps, by addition likewise. 

As a dibasic acid, sulphurous acid forms normal salts like 
Na^SOs, and acid salts like NaHS0 3 . 

Consecutive Reactions. — There are many chemical reactions 
that proceed in two stages, which can be carried out separately. 
This is the case with the two reactions used in the chamber process 
(p. 281). The actions are consecutive, because the second uses 
materials produced by the first. It may be noted that if the 
second action is as speedy as the first, or speedier, then no inter- 
mediate products will be detectable. This is the case with the 
chamber process reactions, when sufficient steam is introduced, 
for under these circumstances no solid nitrosylsulphuric acid is 
deposited. If the second reaction is slower than the first, then the 
products of the first reaction will accumulate, and become notice- 
able. 

The conception of consecutive reactions enables us to under- 
stand and remember some facts. For example, it was mentioned 
that when dry sulphur is oxidized, we obtain sulphur dioxide, but 
when moist sulphur is oxidized, by the air or otherwise, the only 
product is sulphuric acid (p. 267). This change may be conceived 
of as proceeding in two stages: 

S + 2 + H 2 -> H 2 S0 3 , 
2H 2 S0 3 + 2 ->2H 2 S0 4 , 



290 COLLEGE CHEMISTRY 

which would be consecutive reactions. Since oxidation of solid 
sulphur can proceed only on the surface, it is slow. Since the 
sulphurous acid is dissolved, and every molecule of it is accessible 
to the dissolved oxygen, or oxidizing agent, the second action 
should be speedier and consume the product of the first action as 
fast as it is formed. It is, therefore, quite natural that no sul- 
phurous acid should be detectable when water is present. 

Sulphites. — The acid sulphites of the alkali metals, KHS0 3 
and NaHS(>3, when in solution, are acid in reaction, owing to the 
appreciable dissociation of the ion HSC>3~. The sulphites are 
readily decomposed by acids to give free sulphurous acid, and the 
latter partly decomposes, yielding sulphur dioxide (p. 275) . 

Calcium bisulphite solution, Ca(HS0 3 )2, is used to dissolve the 
lignin out of wood, and leave the pure cellulose (paper pulp) 
employed in the manufacture of paper (q.v.) . 

When heated, sulphites undergo decomposition. The sulphates, 
being the most stable of all the salts of sulphur acids, are formed 
when the salts of any of those acids are decomposed by heating. 
The nature of the particular salt determines what other products 
shall appear. Thus, with sodium sulphite Na2S03, one molecule 
of the sulphite furnishes three atoms of oxygen, sufficient to oxi- 
dize three other molecules, and leaves one molecule of sodium 
sulphide behind: 

4Na 2 S0 3 -> Na 2 S + 3Na 2 S0 4 . 

The sulphites are as readily oxidized as is the acid itself. They 
are slowly converted, both in solution and in the solid form, by the 
influence of the oxygen of the air, into sulphates. 

Thiosulphuric Acid H 2 S 2 3 . — This acid is not known in the 
free condition, but its salts are in common use in the laboratory 
and commercially. Sodium thiosulphate, for example, is pre- 
pared by boiling a solution of sodium sulphite with free sulphur. 
The action is something like the addition of oxygen to sulphurous 
acid: 

Na 2 S0 3 + S -> Na 2 S 2 3 or S0 3 = + S -> S 2 3 =. 

Sodium thiosulphate ("hypo") is used in "fixing" photographs. 



THE OXIDES AND OXYGEN ACIDS OF SULPHUR 291 

By the addition of acids to a solution of sodium thiosulphate, 
the thiosulphuric acid is set free, but the latter instantly decom- 
poses, giving a precipitate of sulphur : 

Na^Os + 2HC1 <± H 2 S 2 3 + 2NaCl, 

EAOs *=* Si + H 2 S0 3 ^ H 2 + S0 2 T . 

Per sulphuric Acid H 2 S 2 Oa. — This, like the other acids just 
mentioned, is unstable, and can be kept only in dilute solution. 
Its salts, however, are coming into use for commercial purposes 
and for " reducing" negatives in photography. The salts are 
prepared by electrolyzing sodium-lrydrogen sulphate NaHS04 in 
concentrated solution (Hugh Marshall). The persulphuric acid, 
formed by the union of the negative ions in pairs as they are 
discharged on the anode, 

2HS0 4 ~ + 2©->H 2 S 2 8 ,* 

undergoes double decomposition with the excess of sodium bisul- 
phate, and the less soluble sodium persulphate crystallizes out. 
The other salts are made by double decomposition from this one. 

Compounds of Sulphur and Chlorine. — When chlorine gas 
is passed over heated sulphur, it is absorbed and sulphur mono- 
chloride, a reddish-yellow liquid, boning at 138°, is obtained. The 
molecular weight of this substance, as shown by the density of its 
vapor, indicates that it possesses the formula S 2 C1 2 . When thrown 
into water, it is rapidly hydrolyzed, producing sulphur dioxide and 
sulphur: 2S 2 C1 2 + 2H 2 -> S0 2 + 4HC1 + 3S. 

Sulphur itself dissolves very freely in the monochloride, and the 
solution is employed in vulcanizing rubber. 

Sulphur dioxide and chlorine gases, when exposed to direct sun- 
light, unite to form a liquid known as sulphuryl chloride S0 2 C1 2 . 
Camphor causes the union to take place much more rapidly, owing 
to some catalytic effect. The compound is a colorless liquid, boil- 
ing at 69°. With water it gives sulphuric acid and hydrogen 
chloride : 

S0 2 C1 2 + 2H 2 -» S0 2 (OH) 2 + 2HC1. 

Graphic Formula of Sulphuric Acid. — The actions just 
mentioned give a clue to the constitution of sulphuric acid. Since 
* See footnote, p. 276. 



292 COLLEGE CHEMISTRY 

chlorine does not combine directly with oxygen, but does com- 
bine readily with sulphur, we may assume that, in the formation 
of sulphuryl chloride, S0 2 + Cl 2 — > S0 2 C1 2 , the chlorine unites 
more intimately with the sulphur in the molecule S0 2 : 

°^ °^ / C1 

cr cr N ci 

The action of water upon the product is presumably similar to 
that of water on phosphorus tribromide (p. 197) : 



0-H O vV O-H 

-*2HCl-f JS( 



o %B >ca h 

/y >C1 H4-0- 



The last is called the structural formula of sulphuric acid. It is not 
thereby implied that the atoms in its molecules are attached pre- 
cisely in this manner, however, but rather that the chemical be- 
havior of the substance, as being partly an oxide and partly an 
hydroxide of sulphur, is symbolized in this fashion. Such graphic 
formulae are of great value in expressing the chemical behavior of 
the complex compounds of carbon. 

Exercises. — 1. What ground is there for assigning the formula 
S0 2 instead of S2O4 to sulphur dioxide (p. 277)? 

2. Explain why nitric acid is completely displaced by the action 
of sulphuric acid on sodium nitrate (p. 282). 

3. What are the relative volumes, (a) of sulphur dioxide and 
nitrogen (p. 150) resulting from the roasting of pyrite (p. 275), (b) 
of air and sulphur dioxide in making sulphuric acid, (c) of nitrogen 
(left) to sulphur dioxide (used) in making sulphuric acid, when 
pyrite is the source? 

4. Make a list of, and classify, the various applications of 
sulphuric acid to the liberation of other acids. 

5. Formulate the behavior of the hydrosulphate-ion (p. 287) 
when a solution of barium chloride is added to a rather concen- 
trated solution of sulphuric acid. 

6. Assign to the proper class of ionic actions (pp. 259, 270), (a) 
the action of iodine on sulphurous acid (p. 289), (6) of sulphur on 
sodium sulphite (p. 290), (c) the formation of persulphuric acid 
(p. 291). 



CHAPTER XXII 

SELENIUM AND TELLURIUM 
THE CLASSIFICATION OF THE ELEMENTS 

Along with sulphur, chemists group two other elements, sele- 
nium (Se, at. wt. 79.2) and tellurium (Te, at. wt. 127.5). If, while 
this and the next page are read, the nature of the chief compounds 
of sulphur is kept in mind, the analogy between the nature and 
chemical behavior of the three elements and their corresponding 
compounds will be obvious (see Chemical relations of the sulphur 
family, below). 

Occurrence and Properties of Selenium Se, — Selenium 
(Gk., the moon) occurs free in some specimens of native sulphur, 
and in combination often takes the place of a small part of the 
sulphur in pyrite (FeS 2 ). It is found free in the dust-flues of the 
pyrite-burners of sulphuric acid works. The familiar forms are, 
the red precipitated variety, which is amorphous and soluble in 
carbon disulphide, and the lead-gray, semi-metallic variety, ob- 
tained by slow cooling of melted selenium, which is insoluble, and 
melts at 217°. In the latter form it has some capacity for con- 
ducting electricity, which is greatly increased by exposure to light 
in proportion to the intensity of the illumination. A photometer, 
using this property, has been devised by Joel Stebbins (1914), for 
measuring the relative intensity of the light of different stars. 
Selenium boils at 680°, and at high temperatures has a vapor 
density corresponding to the formula Se2. 

The element combines directly with many metals, burns in 
oxygen to form selenium dioxide, and unites vigorously with 
chlorine. 

Compounds of Selenium. — Ferrous selenide, made by heat- 
ing iron filings with selenium (c/. p. 13), when treated with con- 
centrated hydrochloric acid gives hydrogen selenide: 

FeSe + 2HC1 U H 2 Set + FeCl* 
293 



294 COLLEGE CHEMISTRY 

The compound is a poisonous gas, which possesses an odor recalling 
rotten horse-radish, and is soluble in water. The solution is faintly 
acid in reaction, and deposits selenium when exposed to the action 
of the air (cf. p. 270). Other selenides, which, with the exception 
of those of potassium and sodium, are insoluble in water, may be 
precipitated by leading the gas into solutions of soluble salts of 
appropriate metals (cf. p. 273). 

The dioxide Se0 2 is a solid body formed by burning selenium. 
Selenious acid H 2 Se0 3 may be made by dissolving the dioxide in hot 
water, or by oxidizing selenium with boiling nitric acid. Unlike 
sulphur (p. 267), the element gives little of the higher acid H 2 Se0 4 
by this treatment. The acid is reduced by sulphurous acid to 
selenium: H 2 Se0 3 + 2H 2 S0 3 -> 2H 2 S0 4 + H 2 + Se. 

No trioxide is known. Selenic acid H 2 Se04, a white solid, is 
made in solution by oxidizing silver selenite with bromine-water 
(which contains hypobromous acid, cf. p. 161), and filtering: 

Br 2 + H 2 <=> HBr + HBrO. 
Ag 2 Se0 3 + HBrO -> Ag 2 Se0 4 + HBr. 
2HBr + Ag 2 SeQ 4 -> 2AgBr| + H 2 SeQ 4 . 
Br 2 + H 2 + A&SeOs -> 2AgBr j + H 2 Se0 4 . 

It is itself a powerful oxidizing agent and, even in dilute solution, 
liberates chlorine from hydrochloric acid: H 2 Se0 4 + 2HC1 — » 
H 2 SeO s + H 2 + Cl 2 . Sulphuric acid (cf. p. 286), on the other 
hand, is an oxidizing agent only in somewhat concentrated form, 
and even then it can oxidize hydrobromic acid (p. 196), but not 
hydrochloric acid. 

Tellurium Te. — Tellurium (Lat., the earth) occurs in sylvan- 
ite in combination with gold and silver. It is a white, metallic, 
crystalline substance, melting at 452° (b.-p. 1400°). The free 
element unites with metals directly, and burns in air to form the 
dioxide. 

The compounds of tellurium are similar in composition and mode 
of preparation to those of selenium. Some differences in chemical 
behavior are significant, however. Thus, tellurious acid H 2 TeOg is 
a very feeble acid and is also somewhat basic, a sulphate (2Te0 2 , 
S0 3 ) and a nitrate (Te 2 3 (OH)N0 3 ) being known. In this respect 



THE PERIODIC SYSTEM 295 

it differs markedly from sulphurous acid. Telluric acid does not 
affect indicators, and is therefore actually more feebly acidic than 
is hydrogen sulphide. Tellurium tetrachloride TeCLi, although 
hydrolyzed by water, exists in solution with excess of hydrogen 
chloride: TeCLt + 3H 2 +± H 2 Te0 3 + 4HC1, showing the telluri- 
ous acid to be basic in properties and the element tellurium to be, 
to a certain degree, a metallic element. 

The Chemical Relations of the Sulphur Family. — It will 
be seen that sulphur, selenium, and tellurium are bivalent elements 
when combined with hydrogen or metals. In combination with 
oxygen they form unsaturated compounds of the form X IV 02, while 
their highest valence is found in SO3, Te03, and EfeSeO^ where 
they must be sexivalent. The general behavior of corresponding 
compounds is very similar. At the same time, there is in all cases 
a progressive change as we proceed from sulphur through selenium 
to tellurium. The elementary substances themselves, for example, 
become more like metals, physically, and they show higher and 
higher melting-points. The affinity for hydrogen decreases, as is 
shown by the increasing ease with which the compounds H 2 X are 
oxidized in air. The affinity for oxygen likewise decreases, for the 
elements become increasingly difficult to raise to the highest state 
of oxidation. On the other hand, the tendency to form higher 
chlorides becomes greater. We note also that the compounds 
H2XO4 become less and less active as acids, and that a basic 
tendency begins to assert itself. 

The Periodic System 

Classification, or the arrangement of facts on the basis of like- 
ness, is part of the method of science. It is needed to make 
possible the systematic description of the ascertained facts, which 
is a great aid to the memory, and to furnish a guide in investigation, 
by suggesting relations and so pointing out directions in which new 
facts of interest may be found. Thus, as an aid to memory, we 
have treated the halogens as one family and S, Se, and Te as an- 
other. In each case, we have presented the properties common to 
all members of the group, and have then pointed out the differ- 
ences. Again, in investigation, as soon as we have discovered that 



296 COLLEGE CHEMISTRY 

sulphur and selenium are allied elements, we realize the direction 
in which fruitful results may be expected, and we proceed to make 
the corresponding compounds and to note the resemblances and 
differences in the conditions for preparation and in the properties 
of the compounds obtained. 

Metallic and Non-Metallic Elements, — Thus far we have 
found the division into metallic and non-metallic elements very 
serviceable for classification in terms of chemical relations (p. 163). 
This distinction we shall continue to employ. The metallic, or 
positive elements (p. 94), (1) form positive radicals and ions con- 
taining no other element (c/. p. 247). Thus the metals give sul- 
phates, nitrates, carbonates, and other salts, which furnish a 
metallic ion, such as Na + or K+, together with the ions S0 4 = , N0 3 ~, 
and C0 3 =. (2) Their hydroxides, KOH, Ca(OH) 2 , etc., give the 
same metallic ion, and the rest of the molecule forms hydroxide-ion. 
That is to say, their hydroxides are bases and their oxides are 
basic. The metallic elements often enter, but only with other 
elements, into the composition of a negative ion, as is the case with 
manganese in K.Mn04, with chromium in K 2 .Cr 2 07, and with silver 
in K.Ag(CN) 2 . 

The non-metallic or negative elements (1) are found chiefly in 
negative radicals and ions. They form no nitrates, sulphates, car- 
bonates, etc., for they could not do so without themselves alone 
constituting the positive ion. We have no such salts of sulphur, 
carbon, or phosphorus, for example. (2) Their hydroxides, al- 
though their formulae may be written C10 2 OH, P(OH) 3 , S0 2 (OH)2, 
furnish no hydroxyl ions, as this would involve the same conse- 
quence. These hydroxides are divided by dissociation, in fact, so 
that the non-metal forms part of a compound negative radical, and 
the other ion is hydrogen-ion, C10 3 .H, P0 3 H.H 2 , S0 4 .H 2 . Their 
oxides are acidic. (3) Their halogen compounds, like PBr 3 (p. 197) 
and S 2 C1 2 (p. 291), are completely hydrolyzed by water, and the 
actions are not, in general, reversible. The halides of the typical 
metals are not hydrolyzed (see Chap. XXXIII), and with those that 
are not typical, the action is reversible. 

The distinction is not perfectly sharp, however. Thus, zinc 
gives both salts like the sulphate, Zn.S0 4 , and chloride, Zn.Cl 2 , and 
compounds like sodium zincate (p. 56), Zn0 2 .Na 2 . 



THE PERIODIC SYSTEM 297 

Classification by Atomic Weights. — Newlands (1863-4) dis- 
covered a surprising regularity that became apparent when the ele- 
ments were placed in the order of ascending atomic weight. 
Omitting hydrogen (at. wt. 1) the first seven were: lithium (7), 
glucinum (9), boron (11), carbon (12), nitrogen (14), oxygen (16), 
fluorine (19). These are all of totally different classes, and include 
first a metal forming a strongly basic hydroxide, then a metallic 
element of the less active sort, then five non-metals of increasingly 
negative character, the last being the most active non-metal 
known. The next element after fluorine (19) was sodium (23), 
which brings us back sharply to the elements that form strongly 
basic hydroxides. Omitting none, the next seven elements were 
sodium (23), magnesium (24.4), aluminium (27), silicon (28.4), 
phosphorus (31), sulphur (32), chlorine (35.5). In this series there 
are three metals of diminishing positiveness, followed by four non- 
metals of increasing negative activity, the last being a halogen very 
like fluorine. On account of the fact that each element resembles 
most closely the eighth element beyond or before it in the list, the 
relation was called the law of octaves. After chlorine the octaves 
become less easy to trace. 

That this periodicity in chemical nature is more than a coinci- 
dence is shown by the fact that the valence and even the physical 
properties, such as the specific gravity, show a similar fluctuation 
in each series. In the first two series the compounds with other 
elements are of the types: 

iid, oia, bo,, ecu, £g- 0H2jFH . 

NaCl, MgCl 2 , MCh, SiCk, jg£' |°£ gjg 7 " 

Thus the valence towards chlorine or hydrogen ascends to four and 
then reverts to one in each octave. The highest valence, shown in 
oxygen compounds, ascends from lithium to nitrogen with values 
one to five, and then fails because compounds are lacking. In the 
second octave, however, it goes up continuously from one to 
seven. 

Again, the specific gravities of the elements in the second series, 
using the data for red phosphorus and liquid chlorine, are: 

Na 0.97, Mg 1.75, Al 2.67, Si 2.49, P 2.14, S 2.06, CI 1.33. 



298 COLLEGE CHEMISTRY 

Mendelejeff's Scheme. — In 1869 Mendelejeff published an 
important contribution towards adjusting the difficulty which the 
elements following chlorine presented, and developed the whole 
conception so completely that the resulting system of classification 
has been connected with his name ever since. Almost simultane- 
ously Lothar Meyer made similar suggestions, but did not urge 
them with the same conviction or elaborate them so fully. The 
table on the following page, in which the atomic weights are ex- 
pressed in round numbers, is a modification of one of Mendelejeff's. 

The chief change from the arrangement in simple octaves is that 
the third series, beginning with potassium, is made to furnish 
material for two octaves, potassium to manganese and copper to 
bromine, and is called a long series. The valences fall in with this 
plan fairly well. Copper, while usually bivalent, forms also a 
series of compounds in which it is univalent. Iron, cobalt, and 
nickel cannot be accommodated in either octave, as their valences 
are always two or three. At the time Mendelejeff made the table, 
three places in the third long series had to be left blank, as a tri- 
valent element [Sc] was lacking in the first octave of the series, and 
a trivalent [Ga] and a quadrivalent one [Ge] in the second. These 
places have since been filled, as we shall presently see. The first 
two (the short) series have been split in the table, as lithium and 
sodium closely resemble potassium, while the remaining members 
of these series fall more naturally over the corresponding elements 
of the second octave of the third series. 

The fourth series (long) is nearly complete. It begins with an 
active alkali metal, rubidium, and ends with iodine, a halogen. 
The rule of valence is strictly preserved throughout the series, and 
in general the elements fall below those which they most closely 
resemble. 

The fifth, sixth, and seventh (long) series are incomplete, but the 
order of the atomic weights and the valence enable us satisfac- 
torily to place all but about ten rare elements. The chemical 
relations to elements of the fourth series justify the position as- 
signed to each. Caesium, for example, is the most active of the 
alkali metals; barium has always been classed with strontium, and 
bismuth with antimony. 

In two cases a slight displacement of the order according to 
atomic weights is necessary. Cobalt is put before nickel because it 



THE PERIODIC SYSTEM 



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300 COLLEGE CHEMISTRY 

resembles iron more closely. Tellurium and iodine are placed in 
that order to bring them into the sulphur and halogen groups 
respectively. Their valence and other chemical relations both 
require this. The general agreement, however, is very remarkable. 

General Relations in the System. — In every octave the 
valence towards oxygen ascends from one to seven, while that 
towards hydrogen, in the cases of the last four elements (when they 
combine with hydrogen at all), descends from four to one. The 
physical properties fluctuate within the limits of each series in a 
similar way. The values of each physical constant for correspond- 
ing members of the successive series do not exactly coincide, how- 
ever. A progressive change, as we descend each vertical column, 
is the rule. Thus the specific gravities (water = 1) of the alkali 
metals rise from lithium (0.53) to caesium (1.87). In the same 
group the melting-points descend from lithium (186°) to caesium 
(26.5°). 

As yet no exact mathematical (quantitative) relation between 
the values for any property and the values of the atomic weights 
has been discovered; only a general (qualitative) relationship can 
be traced. Anticipating the discovery of some more exact mode 
of stating the relationship in each case, and remembering that 
similar values of each property recur periodically, usually at inter- 
vals corresponding to the length of an octave or series, the principle 
which is assumed to underlie the whole, the periodic law, is stated 
thus: All the properties of the elements are periodic functions of 
their atomic weights. 

That the chemical relations of the elements vary just as do the 
physical properties of the simple substances is easily shown. 
Thus, each series begins with an active metallic (positive) element, 
and ends with an active non-metallic (negative) element, the inter- 
vening elements showing a more or less continuous variation 
between these limits. Again, the elements at the top are the least 
metallic of their respective columns. As we descend, the members 
of each group are more markedly metallic (in the first columns), or, 
what is the same thing, less markedly non-metallic (in the later 
columns; cf. p. 296). 

In the first series boron is the first non-metal we encounter. In 
the second series silicon is the first such element. In the third 






THE PERIODIC SYSTEM 301 

there is more difficulty in deciding. Titanium, vanadium, and 
germanium are usually, though with questionable propriety, 
classed as metallic elements.* Selenium is undoubtedly non- 
metallic. Arsenic is, on the whole, non-metallic. In the fourth 
series tellurium is commonly considered to be the first non-metallic 
element. Thus a zigzag line, shown in the table, separates all the 
non-metallic elements from the rest of the elements, and confines 
them in the right-hand upper corner. 

A more compact form of the table is printed at the end of this 
book, opposite the rear cover. The only difference between this 
and the other is that the two octaves of each long series have been 
placed in the same set of seven main columns. The iron, palla- 
dium, and platinum groups occupy a column on the right of the 
main columns, and are often called collectively the eighth group. 
The newly discovered elements, found chiefly in the air, have been 
placed at the left-hand side. Since they do not enter into com- 
bination at all, their valence may appropriately be given as zero. 
With the exception of argon, the values of their atomic weights 
agree well with this assignment. Hydrogen is the only common 
element whose place is still in debate. The valence is shown by 
the general formulae at the head of each column. 

Applications of the Periodic System. — The system has 
found application chiefly in four ways: 

1. In the prediction of new elements. Mendelejeff (1871) drew 
attention to the blank then existing between calcium (40) and tita- 
nium (48). He predicted that an element to fit this place would 
have an atomic weight 44 and would be trivalent. From the 
nature of the surrounding elements, he very cleverly deduced many 
of the physical and chemical properties of the unknown element and 
of its compounds. In 1879 Nilson discovered scandium (44), and 
its behavior corresponded closely with that predicted. Mendele- 
jeff described accurately two other elements, likewise unknown at 
the time. In 1875 Lecoque de Boisbaudran found gallium, and 
in 1888 Winkler discovered germanium, and these blanks were 
filled. 

* In discussing chemical relations, the term metallic element is preferable 
to metal. The free element (e.g., arsenic) may have the luster of a metal, 
and yet the element, in combination, may be non-metallic or negative. 



302 COLLEGE CHEMISTRY 

2. By enabling us to decide on the correct values for the atomic 
weights of some elements, when the equivalent weights have been 
measured, but no volatile compound is known (cf. pp. 104 and 118). 
Thus, the equivalent weight of indium was 38 and, as the element 
was supposed to be bivalent, it received the atomic weight 76. It 
was quite out of place near arsenic (75), however, being decidedly 
a metallic element. As a trivalent element with the atomic weight 
115, it fell between cadmium and tin. Later work fully justified 
the change. Quite recently, radium has been discovered, and 
found to have the equivalent weight 113 and to resemble barium. 
If, like barium, it is bivalent, it occupies a place under this ele- 
ment, in the last series. 

3. By suggesting problems for investigation. The periodic system 
has been of constant service in the course of inorganic research, and 
has often furnished the original stimulus to such work as well. For 
example, the atomic weight of tellurium bore the value 128 when 
the table was first constructed, and it was confidently expected that 
reexamination would bring this value below that of iodine (then 
127, now 126.92). Several most careful studies of the subject have 
been made by different methods. It seems probable that the real 
value of the atomic weight is not far from Te = 127.5, and there- 
fore more than half a unit greater than that of iodine. Since, how- 
ever, mathematical correspondence is found nowhere in the system, 
the existence of marked inconsistencies like this need not shake our 
confidence in its value when it is used with due consideration of the 
degree of correspondence to be expected. 

In the same way, incorrect values of many physical properties 
have been detected, and have been rectified by more careful 
work. 

4. By furnishing a comprehensive classification of the elements, 
arranging them so as to exhibit the relationships among the physical 
and chemical properties of the elements themselves and of their 
compounds. Constant use will be made of this property of the 
table in the succeeding chapters. Having disposed of the halogen 
and sulphur families (excepting the oxygen compounds of the 
former), situated, respectively, in the seventh and sixth columns 
of the table (at the end of this book), we shall presently take up 
nitrogen and phosphorus from the right side of the fifth column. 
Then from the fourth column, we shall select carbon and silicon, 



THE PERIODIC SYSTEM 303 

and from the third boron, leaving the other, more decidedly metal- 
lic elements for later treatment. 

Moseley's Atomic Numbers. — We have seen that simple, 
mathematical relations between the atomic weights and the physi- 
cal or chemical properties of an element do not exist. In several 
instances, the atomic weights are not even in the same order as 
are the values of the properties. We have now obtained from 
another direction numbers which seem to be more fundamental 
even than atomic weights. 

Visible light, X-rays, and wireless electric waves are all vibra- 
tions of the same nature in the ether. They differ only in wave- 
length, the order of the wave-lengths being 10~ 5 cm., 10~ 8 cm., and 
10 6 cm. (10 kilometers), respectively. Now, just as the spectrum 
of visible light is obtained by using a grating, on which the rulings 
are separated by distances of the order of the wave-length of such 
light, so ordinary crystals give spectra of X-rays, because they are 
composed of particles arranged in rows about one thousand times 
closer and so form a suitable grating for X-rays. This fact was 
first discovered by Dr. Laue of the University of Zurich (1912). 
The X-rays are produced in 
an evacuated tube by cathode 
rays, which are streams of 
electrons emanating from 
the cathode (C, Fig. 88), 
when they strike the anti- 
cathode (A). 

With different elements on the anti-cathode, X-rays of slightly 
different wave-lengths, and therefore giving different X-ray 
spectra, are produced. By using different elements, Moseley 
(1914) has found that the higher the atomic weight the shorter the 
wave-length of the characteristic X-rays. When the elements 
are arranged in the order of these wave-lengths, whole numbers 
can be assigned to each which are inversely proportional to the 
square roots of the wave-lengths of corresponding lines in their 
X-ray spectra. These atomic numbers have been determined for 
most of the elements, the atomic weights of which lie between 
those of aluminium and gold. In the following table, the atomic 
numbers for these elements are given and numbers for the twelve 
elements preceding Al have been inserted also. 




304 



COLLEGE CHEMISTRY 



ATOMIC NUMBERS (Moseley) 



He 2 


Li 3 


Gl 4 


B 5 


C 6 


N 7 


8 


F 9 






Ne 10 


Na 11 


Mg 12 


All 13 


Si 14 


P 15 


S 16 


CI 17 






A 18 


K 19 


Ca 20 


Sc 21 


Ti 22 


V 23 


Cr 24 


Mn25 


Fe 26 


Co 27 




Cu 29 


Zn 30 


Ga 31 


Ge 32 


As 33 


Se 34 


Br 35 






Kr 36 


Rb37 


Sr 38 


Y 39 


Zr 40 


Cb 41 


Mo 42 


— 43 


Ru44 


Rh45 




Ag 47 


Cd 48 


In 49 


Sn 50 


Sb 51 


Te 52 


I 53 






Xe 54 


Cs 55 
Au 79 


Ba 56 


La 57 


Ce 58 


Ta 73* 


W 74 


— 75 


Os 76 


Ir 77 



Ni 28 
Pd 46 
Pt 78 



* The atomic numbers 59-72 are those of the metals of the rare earths: Pr 59, Nd 60, 
-61, Sa 62, Eu 63, Gd 64, Tb 65, Dy 66, Ho 67, Er 68, Tm 69, Yb 70, Lu 71, -72. 

It will be seen that there is a whole number available for every 
known element, up to and including gold, and not omitting the 
rare elements which have no satisfactory place in the periodic 
system. There are two blank numbers in the table, which corre- 
spond to two spaces below Mn in the periodic system, and two 
more amongst the rare elements, indicating only four elements 
with atomic weights less than that of gold yet to be discovered. 
The atomic numbers of argon and potassium place them in the 
chemically correct order, while the atomic weights do not. The 
same is true of cobalt and nickel and of tellurium and iodine. 
Finally, it is evident that the atomic weight of each element is, 
roughly, double its atomic number. 

The atomic numbers represent the number of unit positive 
charges of electricity in the nucleus of the atom of each element 
(p. 235). Rutherford has shown that the nucleus contains almost 
the whole mass of the atom, although one or more electrons 
(negative) are present also. Thus, the positive nucleus of the 
hydrogen atom is 1800 times heavier than one electron. The 
nucleus, however, is very minute, having a diameter only about 
one-eighteen hundredth of that of an electron. 

The atomic numbers apparently determine all the properties 
of each element, and are more fundamental than the atomic 
weights. The latter are secondary properties, in most cases 
modified by other factors, and in a few cases actually thrown out 
of order by such factors. 

Crystal Structure. — In this connection it may be mentioned 
that by using crystals of different substances as X-ray gratings, 



THE PERIODIC SYSTEM 305 

W. L. Bragg (1914) has been able to measure the distances be- 
tween the rows of particles in crystals. He also finds that the 
particles, the regular arrangement of which gives the structure 
(p. 82) of the crystal (e.g.-, a cube of common salt), are not the 
molecules of the compound, much less aggregates of such mole- 
cules, but the atoms of the constituent elements. It would thus 
appear that the physical forces (if we may call them physical) 
which hold the crystalline solid together have completely crushed 
the chemical, molecular structure out of existence, and have ar- 
ranged the constituent atoms, as the units of the structure, in a 
crystallographic pattern. Of course, when the crystal-form is de- 
stroyed, by melting, solution, or vaporization, the neighboring 
atoms remain united in groups, constituting the chemical mole- 
cules of the substance. 

Exercises, — 1. Can you explain the presence of free selenium 
in the flues of pyrite burners (p. 294)? 

2. How should you attempt to obtain H 2 Te, and what physical 
and chemical properties should you expect it to possess? 

3. Make a list of bivalent elements and criticize this method of 
grouping as a means of chemical classification. 

4. Write down the symbols of the elements in the fourth series 
(that beginning with rubidium, and ending with iodine) on p. 299. 
Record the valence of each element toward oxygen, using for refer- 
ence the chapters in which the oxygen compounds are described. 



CHAPTER XXIII 

OXIDES AND OXYGEN ACIDS OF THE HALOGENS 
OXIDATION AND REDUCTION 

The chief subjects of practical importance touched upon in the 
first part of this chapter are connected with bleaching powder 
CaCl(OCl), and potassium chlorate KC10 3 and perchlorate KC10 4 . 
Hence our attention will be largely directed to the modes of making 
these substances and to their relations to one another. Inciden- 
tally, we shall encounter many actions of a complex and, to us, 
more or less novel kind. 

Compounds of Chlorine Containing Oxygen. — The fol- 
lowing are the names and formulae of the substances: 

HCIO Hypochlorous acid, C1 2 Hypochlorous anhydride, 

[HCIOJ Chlorous acid, 

CIO2 Chlorine dioxide, 

HCIO3 Chloric acid, 

HCIO4 Perchloric acid, CI2O7 Perchloric anhydride. 

There are also salts of these acids, like the three substances 
mentioned in the first paragraph. Chlorous acid is itself unknown, 
but potassium chlorite KCIO2 and some other derivatives have 
been made. 

The two anhydrides (p. 94), when brought into contact with 
water, combine with it to form the acids opposite which they stand 
in the table. Chlorine dioxide (q.v.) } however, is not related to any 
one acid in this way. 

All these compounds differ from most that we have hitherto dis- 
cussed, inasmuch as not one of them can be made by direct union 
of the simple substances. 

Nomenclature of the Acids and their Salts. — The acids 
and salts are named on a plan similar to that used in the case of 
the sulphur acids: 

306 



OXIDES AND OXYGEN ACIDS OF CHLORINE 307 

KCIO Potassium hypochlorite, HGO Hypochiorous acid, 

KC10 2 Potassium chlorite, HC10 2 Chlorous acid, 

KCIO3 Potassium chlorate, HCIO3 Chloric acid, 

KCIO4 Potassium perchlorate. HC10 4 Perchloric acid. 

It should be noted, however, that the use of ic and ous for more and 
less oxygen, respectively, and of hypo for still less and of per for 
still more oxygen are simply relative terms within a single group. 
Thus, sulphuric acid H 2 S0 4 has a composition entirely different 
from chloric acid, and both of these differ in composition from 
phosphoric acid H3PO4. The names and formulae of each group 
must be learned, separately. 

Chlorine Monoxide or Hypochlorous Anhydride ChO. — 

A solution of pure hypochlorous acid is most easily prepared by 
dissolving the anhydride in water. This oxide is obtained by 
passing chlorine gas over warmed mercuric oxide * HgO (Fig. 66, 
p. 156). Each of the constituents of the oxide combines with 
chlorine: 

HgO + 2C1 2 -» HgCl 2 + C1 2 0. 

The mercuric chloride then unites with another formula-weight of 
the mercuric oxide to form a solid basic mercuric chloride HgO, 
HgCl 2 , which remains in the tube. The chlorine monoxide is a 
brownish-yellow, heavy, easily liquefied gas (b.-p. 5°). When 
slightly warmed it decomposes into its constituents with explosion. 
The gas dissolves in water very easily (200 : 1, by vol.). The yel- 
low solution of hypochlorous acid which results: 

C1 2 + H 2 t? 2HOC1, 

has a strong odor of chlorine monoxide, because the combination 
is reversible. There are other ways of preparing a dilute solution 
of the acid (see below). 

Properties of Hypochlorous Acid, — Hypochlorous acid is 
unstable, and cannot be made, excepting in solution, or kept, ex- 

* The crystalline, red oxide is not sufficiently active. The oxide must be 
precipitated from sodium hydroxide and mercuric nitrate solutions, it must 
be washed thoroughly on a filter, and be dried at 300-400° before use. 



308 COLLEGE CHEMISTRY 

cepting in dilute solution. This is in consequence of its tendency 
to decompose in three different ways, one of which, the liberation 
of the anhydride, has just been mentioned. 

1. Hypochlorous acid is a little-ionized, weak add. 

HOC1 <± H+ + CIO". 

It neutralizes active bases, its ionization equilibrium being dis- 
placed forwards as the hydrogen-ion H + is removed to form water: 

NaOH + HOC1 *± NaOCl + H 2 0. 

2. The solution, if strong, gives off chlorine monoxide C1 2 0, the 
union with water being reversible. 

3. If the solution is concentrated, much of the hypochlorous 
acid changes gradually into chloric acid and hydrogen chloride. 
This is a self -oxidation. It occurs even in the dark: 

3HOC1 -> HC10 3 + 2HC1. 

4. When the solution is exposed to sunlight, oxygen is evolved 
rapidly. 

2HOCl->2HCl + 2 . 

This decomposition always takes place in sunlight, whether the 
acid is present alone in the water, or along with other substances. 
We have already noted this fact in discussing chlorine-water (p. 
162), which contains this acid. 

5. In consequence of the ease with which it gives up oxygen, 
hypochlorous acid is a strong oxidizing agent. In this direction 
it has several important commercial applications (see below). 

Commercial Preparation of Hypochlorites. — For com- 
mercial purposes, pure hypochlorites are not, as a rule, required. 
Hence, sodium or potassium hypochlorite is prepared by the action 
of sodium or potassium hydroxide on chlorine-water. The latter 
contains both hydrochloric and hypochlorous acids, and so a 
solution containing a mixture of sodium or potassium chloride 
and hypochlorite is obtained: 

Cl 2 + H 2 <± HC1 + HOC1. (1) 

HC1 + KOH <=± KC1 + H 2 0. (2) 

HOC1 + KOH *± KOC1 + H 2 0. (3) 



OXIDES AND OXYGEN ACIDS OF CHLORINE 309 

Although action (1) is only partial, being strongly reversible, the 
neutralization of the two acids in actions (2) and (3) displaces the 
first equilibrium-, and all three actions proceed to completion. 
Action (1), followed by (2) or (3), is a pair of consecutive actions 
(p. 289), of which the second (the neutralization) is the speedier 
of the two. Both pairs of consecutive actions (1) + (2) and (1) 
+ (3), can be combined in one equation. Thus, omitting the 
water, which appears both among products and initial substances 
and in any case is present in large excess as a solvent, and omitting 
also the two acids, which are used up as quickly as they are pro- 
duced by equation (1) and are not amongst the actual products, 
we get, by addition of the three equations (cf. p. 195), the final 
equation: 

Cl 2 + 2KOH -» KC1 + KOC1 + H 2 0. 

As lime is a less expensive alkali than is potassium or sodium 
hydroxide, it is largely used. The chlorine is led into chambers 
-ontaining quicklime CaO spread on trays: 

/ C1 
N OCl 

The product is not a mixture, but a mixed salt (p. 245), known 
as bleaching powder or "chloride of lime." The fact that this 
is a mixed salt does not interfere with its use as a commercial 
source of hypochlorous acid. It is only moderately soluble in 
water. 

Hypochlorous Acid from Bleaching Powder, — 1. When 
bleaching powder is dissolved in water, being a salt, it is very ex- 
tensively ionized (see formulation) . If now an active acid, that is, 
one giving a large concentration of hydrogen-ion, is added, the 
values of the products of the concentrations [H + ] X [Cl~] and 
[H + ] X [OCl~], on which depend the extent to which molecules of 
HC1 and HOC1 will be formed (p. 238), are large. HCIO, being 
little ionized, is formed extensively: HC1, being highly ionized is 
formed in much smaller amount. Both, however, interact to pro- 
duce chlorine and water, and this displaces the other equilibria. 
Hence an active acid decomposes the salt almost completely. An 



310 COLLEGE CHEMISTRY 

active acid gives, therefore, chlorine-water, and not pure hypo- 
chlorous acid. 

CaCl(OCl) t± Ca++ + CI" + OC1" 2 * A weak acid > however, 

H 2 SQ 4 <=*S0 4 = + H+ +H+ llke bonc acid or carbonic 

it if acid, gives so low a concen- 

HC1 HOC1 ^ ra ti° n of H+ that union of 

' 7? this ion with OC1" occurs to 

$ form the little ionized HOC1 

2 ~*~ 2 only, and practically no com- 
bination of H+ with Cl~ takes place (see bleaching). 

CaCl(OCl) <± Ca++ + CI" + OCl"1 _> nnni 
H 2 C0 3 ^±C0 3 = + H+ + H+ f^^OU. 

When the dilute mixture is distilled, chlorine monoxide (2HOC1 <=* 
H 2 + C1 2 0) passes over with the steam, and so a dilute hypo- 
chlorous acid can be obtained. 

Hypochlorous Acid from Chlorine-Water. — An interesting 
way of obtaining dilute hypochlorous acid is to add chalk CaC0 3 
to chlorine-water and distil. Here, the chalk is insoluble, and so 
gives a very low concentration of Ca++ + C0 3 ~. The HC1 in the 
chlorine-water gives, however, a sufficiently large concentration 
of H + to combine with the C0 3 = to form H 2 C0 3 , which is hardly 
ionized at all. This carbonic acid H 2 C0 3 then decomposes and 
carbon dioxide is liberated: 

CaC0 3 (solid)^CaC0 3 (dslvd)^Ca+++C0 3 =l .rr rn _,„ n , m 
2HC1 <=±2C1"+2H+ pH 2 C0 3 ^H 2 0+C0 2 . 

The hypochlorous acid, however, remains molecular HOC1, gives 
almost no H+, and so for the most part remains unaffected. It 
can afterwards be distilled off with the water. 

Hypochlorous Acid as an Oxidizing Agent, — Hypochlorous 
acid, in decomposing into oxygen and hydrochloric acid, gives 
off heat. HOCl,Aq -* HCl,Aq + O + 9300 cal. Hence more 
energy is liberated in oxidation by the acid than in oxidation by 
free oxygen, and the former is therefore more active as an oxidizing 
agent (p. 224). Thus, hypochlorous acid, either in pure solution 
or in the form of chlorine-water, oxidizes sulphurous acid instantly: 
H 2 S0 3 + HOC1 -> H 2 S0 4 + HC1. 



OXIDES AND OXYGEN ACIDS OF CHLORINE 311 

It also oxidizes bromine and iodine, in water, although these ele- 
ments are not affected by free oxygen, giving bromic and iodic 
acids, respectively: 

5HC10 + I 2 + H 2 -» 5HC1 + 2HI0 3 . 

The solution also oxidizes organic colored substances (p. 221), 
producing colorless, or less strongly colored ones. Thus, it 
oxidizes indigo (deep blue) quickly to isatin, a yellow substance 
relatively pale in color: 

Ci 6 H 10 N 2 O 2 + 2HOC1 -* 2C 8 H 5 N02 + 2HC1. 

In ways just as definite as this, hypochlorous acid will change the 
composition of other colored substances, although, since we do not 
know the formulae of all these substances, we cannot always write 
equations for the actions. 

Hypochlorous Acid as a Bleaching Agent. — It is on account 
of its oxidizing power that hypochlorous acid is used commercially 
in bleaching. It is not applied to paints, which are chiefly mineral 
substances, but to complex compounds of carbon, such as consti- 
tute the coloring matters of plants and of those artificial dyes 
which are now manufactured in great variety. 

Cotton and linen, in their original states, are not pure white. 
Bleaching is, therefore, an extensive and most important industry. 
The yarn or cloth must first be freed from cotton-wax and tannin, 
since the former would protect it from the action of the bleaching 
agent, and both would make the subsequent dyeing uneven. The 
material is, therefore, first boiled with very dilute sodium hydroxide 
solution, and washed with water. The goods are then saturated 
with bleaching powder solution, and piled loosely until the coloring 
matter has been oxidized. They are finally washed with extreme 
thoroughness. 

As a rule, an active acid is not added. The bleaching is pro- 
duced by the hypochlorous acid liberated by the action of the car- 
bon dioxide from the air. The carbon dioxide dissolves in the 
water of the solution on the goods, and forms carbonic acid: 
C0 2 + H 2 *± H2CO3 (see p. 310, par. 2). The subsequent wash- 
ing removes all traces of the bleaching powder, of the lime which 
the powder often contains, and of the hypochlorous acid, which 



312 COLLEGE CHEMISTRY 

otherwise would act gradually upon the cotton or linen and 
"rot" it. Bleaching agents, when used in the household with- 
out sufficiently careful washing, are liable to cause serious damage 
from this cause. 

Cotton and linen are composed of cellulose (C 6 Hi O 5 ) x , a rather 
inert substance, and one which is very slowly acted upon by dilute 
hypochlorous acid. Hence, with brief contact and proper han- 
dling, no damage is done. Wool, silk, and feathers, however, are 
composed largely of compounds (proteins) containing nitrogen 
(up to 15 per cent) in addition to the above three elements. Their 
constituent material interacts as easily with hypochlorous acid as 
do the traces of coloring substances. Hence, since the fabric itself 
would be attacked by this agent, sulphur dioxide or sulphurous 
acid (p. 289) is used for bleaching these materials. 

Bleaching Powder in Sanitation. — A disinfectant is a sub- 
stance which destroys bacteria and other minute organisms. 
Bleaching powder has a distinct odor of chlorine monoxide (not 
chlorine). This is due to the action of atmospheric carbon di- 
oxide liberating hypochlorous acid (p. 310). The dry powder 
therefore will disinfect the air and surrounding objects. It must 
be used with discretion, however, as the gas is very corrosive. 

As already mentioned (p. 91), in the purification of city waters 
the organisms which give rise to typhoid fever are destroyed by 
adding a small proportion of bleaching powder solution (about 20 
lbs. per million gallons of water). The salt is hydrolyzed (p. 197), 
giving a basic calcium chloride and free hypochlorous acid. The 
latter kills the organisms, and is itself decomposed in the process, 
so that nothing offensive remains in the water. There is only a 
minute increase in the proportion of salts of calcium (hardness). 

Recently, chlorine-water, made by use of cylinders of liquid 
chlorine (p. 160), has in many cases taken the place of bleaching 
powder solution for this purpose. 

Chlorine not a Bleaching Agent, — Chlorine itself is often, 
erroneously, described as a bleaching agent. If a dry, colored 
cloth be hung for a week in chlorine gas, dried by a little sulphuric 
acid in the bottom of the bottle (Fig. 89), little or no change in the 
color will occur. But a wet rag is bleached as soon as the chlorine 



OXIDES AND OXYGEN ACIDS OF CHLORINE 



313 



has time to dissolve in the water and give the necessary hypochlo- 
rous acid. Flowers are bleached by dry chlorine gas, because by 
their nature they contain the indispensable water. 



Chemical Properties of Hypochlorites. — 

When hypochlorites are heated they change into 
chlorates (see below). They may also give off 
oxygen, 2CaCl(OCl) -> 2CaCl 2 + 2 . Although 
this decomposition is slow in cold solutions of 
hypochlorites, or when they are preserved in the 
dry form, it may be hastened by means of cata- 
lytic agents. The addition of a little cobalt hy- 
droxide (q.v.) to a paste of bleaching powder and 
water causes rapid evolution of oxygen. 




Fig. 89. 



Chlorates. — Like hypochlorous acid itself, the hypochlorites 
turn into chlorates. Thus, when chlorine is passed into a warm, 
concentrated solution of potassium hydroxide, and particularly 
tvhen an excess of chlorine is used, the potassium hypochlorite 
changes into potassium chlorate KC10 3 as fast as it is formed. 
Since this action (equation 2) requires 3KC10, the equation 
formerly given (p. 309) must be tripled: 

3C1 2 + 6KOH -> 3KC1 + 3KC10 + 3H 2 0. (1) 

3KC10-»2KC1 +KCIO3. (2) 

Adding: 3C1 2 + 6KOH -> KC10 3 + 5KC1 + 3H 2 0. 

When the solution is cooled, the less soluble chlorate crystallizes. 

This action involves converting five-sixths of the valuable potas- 
sium hydroxide into the relatively less valuable potassium chloride. 
Hence, in practice, the makers carry out the corresponding action 
with calcium hydroxide. They then add potassium chloride to the 
resulting solution, containing calcium chloride (very soluble) and 
calcium chlorate Ca(C10 3 ) 2 . The potassium chlorate, formed 
by double decomposition, crystallizes when the solution is 
cooled. 

All chlorates are at least moderately soluble in water (see Table 
inside of front cover). Potassium chlorate is used in making fire- 
works, explosives, and matches. An intimate mixture with sugar 
C^^Ou burns with semi-explosive violence, the oxygen of the 



314 COLLEGE CHEMISTRY 

salt combining with the carbon and hydrogen of the sugar to form, 
carbon dioxide and water. 

Chloric Acid HC10 3 . — Since none of the acids of this series 
can be obtained by direct union of their elements (p. 306), it is 
usual first to prepare the salts, and to make the acids from the 
salts by double decomposition. This acid may be obtained, in 
solution in water, by adding the calculated amount of diluted 
sulphuric acid to a solution of barium chlorate: 

Ba(C10 3 ) 2 + H2S.O4 *=► BaS04 + 2HC10 3 . 

The barium sulphate, being insoluble, is removed by nitration. 
It will be noted that double decomposition involving precipitation 
may thus be used for obtaining a soluble product, as well as an in- 
soluble one (cf. selenic acid, p. 294). 

The solution may be concentrated (to about 40 per cent) by 
evaporation, but must not be heated above 40°, as the acid decom- 
poses near this temperature. The resulting thick, colorless liquid 
has powerful oxidizing qualities, setting fire to paper (made of 
cellulose (CeHioOs)*) which has been dipped into it. It converts 
iodine into iodic acid, 2HC10 3 + I 2 -> 2HI0 3 + Cl 2 . When 
not in solution, or when warmed in solution beyond 40°, the acid 
decomposes, giving chlorine dioxide and perchloric acid: 

3HC10 3 -> H 2 + 2C10 2 + HC10 4 . 

Chlorine Dioxide: Chlorous Acid. — Chlorine dioxide C10 2 

(see above) is a yellow gas which may be liquefied, and boils at 
+ 10°. The gas and liquid are violently explosive, the substance 
being resolved into its elements with liberation of much heat. It 
is formed whenever chloric acid is set free, and hence it is seen 
when a little powdered potassium chlorate is touched with a drop 
of concentrated sulphuric acid (end of last section)*. Concen- 
trated hydrochloric acid turns yellow from the same cause when 
any chlorate is added to it. These actions are used as tests for 
chlorates, and distinguish them from perchlorates (q.v.). With 

* The mixture of sugar and potassium chlorate (p. 313) can be set on fire 
by a drop of sulphuric acid. The latter liberates chloric acid, which in turn 
gives C10 2 , and the latter, being a violent oxidizing agent, starts the combus- 
tion of the sugar. 



OXIDES AND OXYGEN ACIDS OF CHLORINE 315 

water, chlorine dioxide gives a mixture of chlorous acid HC10 2 
and chloric acid, and with bases a mixture of the chlorite and 
chlorate. 

Perchlorates. — When heated, chlorates give perchlorates. 
Chlorates also give oxygen at the same time (p. 27): 

(2KC10 3 -+2KCl + 30 2 , 
f 4KC10, -> 3KC10 4 + KC1. 

These actions, like the three decompositions of hypochlorous 
acid (p. 308), are independent, and proceed simultaneously. 
They are concurrent reactions (see below). Their relative speed, 
however, varies with the temperature, and the decomposition into 
chloride and oxygen may completely outrun the other when a 
catalytic agent like manganese dioxide is added (p. 29). When 
pure potassium chlorate is heated cautiously, about one-fifth of 
it has lost all its oxygen by the time the rest has turned into per- 
chlorate. The mixture may be separated by grinding with the 
minimum quantity of water which will dissolve the chloride it con- 
tains. The perchlorate, having at 15° less than one-twentieth of 
the solubility of the chloride, will remain, for the most part, un- 
dissolved. The perchlorates are much more stable (p. 93) than 
the chlorates, or hypochlorites: they are all soluble in water, and 
they are used in making matches and fireworks. 

Perchloric Acid HCIO^ and Perchloric Anhydride Cl 2 7 . — 

Pure perchloric acid explodes when heated above 92°. But, like 
other liquids, its boiling-point is lower when its vapor is under 
reduced pressure (cf. p. 87). At 56 mm. pressure it boils at 39°, a 
temperature at which hardly any decomposition is noticeable. 
Hence the acid may be made by mixing potassium perchlorate 
and concentrated sulphuric acid and distilling the mixture cau- 
tiously in a vacuum (p. 222) : 

KC10 4 + H 2 S0 4 *=; KHSO4 + HCIO4T. 

Perchloric acid is a colorless liquid, which decomposes, and often 
explodes spontaneously, when kept. A 70 per cent solution in 
water is perfectly stable, however. Although it is an active oxidiz- 
ing agent, it is not so active as chloric acid, and does not oxidize 



316 COLLEGE CHEMISTRY 

hydrogen chloride in cold aqueous solution. Hence a drop of 
hydrochloric acid placed on a crystal of a perchlorate gives no 
yellow color. When the acid is liberated by concentrated sulphuric 
acid, it does not at once give the yellow chlorine dioxide (p. 314). 
Perchloric anhydride CI2O7 may be prepared by adding phosphoric 
anhydride to perchloric acid in a vessel immersed in a freezing mix- 
ture, P2O5 + 2HCIO4 -> 2HP0 3 + C1 2 7 . Phosphoric anhydride 
is often used in this way for removing the elements of water from 
compounds. It combines with the water to form metaphosphoric 
acid HPO3. By gently warming the mixture, the perchloric 
anhydride can be distilled off. It is a colorless liquid boiling at 
82° (760 mm.) and exploding when struck or too strongly heated. 

Relation of Anhydride and Acid or Salt. — The derivation 
of the formula of the anhydride from that of the acid or salt 
should receive special attention. In the mind of the chemist, the 
one always instantly suggests the other, so often does he think 
of them as potentially the same substance. The beginner, how- 
ever, finds this habit hard to acquire, and indeed is more likely to 
blunder, in trying to divide the formula of an acid into the formulae 
of water and the anhydride, than in any other calculation he makes. 

The rule is: If the formula of the acid shows an even number 
of hydrogen atoms (H 2 S0 4 or EL^SiC^), subtract all the elements 
of water (H 2 or 2H 2 0), and the balance is the anhydride (SO3 
or Si0 2 ). The divided formulae are H 2 0,S0 3 or 2H 2 0,Si0 2 . If 
there is an odd number of hydrogen atoms (HC10 4 or H3PO4) 
double the formula (H 2 C1 2 8 or H 6 P 2 8 ), and subtract all the ele- 
ments of water as before (C1 2 7 or P 2 5 ) . Then check the result, 
by adding the water again, and dividing by two, correcting the 
blunder if one has been made. 

If the substance is a salt (CuS0 4 or KC10 4 ), subtract the oxide 
of the metal (CuO or K 2 0), taking care to assign to the metal the 
same valence in the oxide as it shows in the salt (SO3 or Cl 2 07). 

There are several uses for this art of ascertaining the anhydride 
corresponding to a given salt or acid. One is in the making of 
equations (see p. 325). Another is in finding the valence of 
the non-metal. Thus, in KC10 4 the anhydride is C1 2 7 , and the 
valence of the chlorine is seven. In H3PO4 the anhydride is P2O5 
and the phosphorus quinquivalent. In HP0 3 (metaphosphoric 



OXIDES AND OXYGEN ACIDS OF CHLORINE 317 

acid), the anhydride is again P 2 5 , and the phosphorus is therefore 
in the same state of oxidation — both are phosphoric acids. 

Simultaneous, Independent Chemical Changes in the 
Same Substances. — When two or more reactions go on simul- 
taneously in the same materials, the actions may be consecutive 
(p. 289) or they may be parallel. In the latter case they are called 
concurrent reactions. Thus, hypochlorous acid undergoes three 
different changes: 

2HC10 -> H 2 + C1 2 0. 

3HC10 -> HCIO3 + 2HC1. 

2HC10 -» 2HC1 + 2 . 

Some molecules decompose into water and chlorine monoxide (p. 
308), while others give chloric acid and hydrogen chloride, and still 
others hydrogen chloride and oxygen. Since the same molecule 
cannot undergo more than one of these different changes, it follows 
that the actions are independent of one another. This is shown 
by the fact that in sunlight the third predominates, while in the 
dark it falls far behind the second. Since the relative quantities 
of the products vary, the several simultaneous actions cannot be put 
in the same equation. The fundamental property of an equation is 
to show the constant proportions by weight between every pair of 
substances in it. Hence three separate equations are required in 
the present, and in all similar cases where all the proportions are 
not constant. Thus, again, in the decomposition of potassium 
chlorate by heating (p. 315), it would be misleading and wrong to 
add the two equations together and write, for the whole action: 

2KCIO3 -> KC1 + KCIO4 + 2 . 

This equation would mean that the proportions amongst the prod- 
ucts were always KC1 : KC10 4 : 2 or 74.6 : 138.6 : 32, whereas, 
in fact, the proportions vary with the conditions — the tempera- 
ture used or the presence of a catalyst which hastens one action 
but not the other. 

Consecutive reactions (p. 289), however, like (1) followed by (2) 
on pp. 308, 313, may be combined in one equation, since in them 
all the proportions must necessarily be constant. These equations 
are interlocked, for (2) consumes what (1) produces. 



318 COLLEGE CHEMISTRY 

Oxygen Acids of Bromine. — No oxides of bromine have been 
made, but the acids HBrO (hypobromous acid) and HBr0 3 (bromic 
acid) and their salts are familiar. 

By the action of bromine on dilute, cold potassium hydroxide 
solution, potassium bromide and hypobromite are formed: 

Br 2 + 2KOH -> KBr + KBrO + H 2 0. 

When the solution is heated, the hypobromite turns into potassium 
bromate and bromide. The actions are exact parallels of the cor- 
responding ones for chlorine (pp. 309, 313). 

Aqueous bromic acid HBr0 3 may be made in the same way as 
chloric acid (p. 314), or by the action of chlorine-water on bromine: 

5HC10 + H 2 + Br 2 -> 2HBr0 3 + 5HCL 

The solution is colorless and has powerful oxidizing properties. 
Thus, it converts iodine into iodic acid : 2HBr0 3 + 1 2 — > 2HI0 3 + Br 2 . 
It appears, therefore, that iodine has more affinity for oxygen than 
has bromine. 

The Oxide and Oxygen Acids of Iodine. — The following are 
the familiar acids and their corresponding salts: 

HI0 3 Iodic acid, KI0 3 Potassium iodate, 

[HIO4 Periodic acid], NaI04 Sodium periodate, 

HbIOg Periodic acid, Na 2 H 3 IC>6 Disodium periodate. 

There is one oxide, iodic anhydride I 2 5 . 

Sodium Iodate NaI0 3 is found in Chile saltpeter. It may 
be made, in much the same fashion as are the chlorates and 
bromates (pp. 313, 318), by adding powdered iodine to a hot solu- 
tion of potassium or sodium hydroxide. It is disodium periodate 
Na 2 H 3 IOe, however, which, being least soluble, crystallizes out. 

Iodic Acid HI0 3 is formed by passing chlorine through iodine 
suspended in water. The action is parallel to that of chlorine on 
bromine-water: 

5HC10 + H 2 + I2 -> 2HI0 3 + 5HC1. 

A still better way is to boil iodine with aqueous nitric acid (q.v.). 
The latter gives up oxygen readily, and is here used solely on this 



OXIDES AND OXYGEN ACIDS, THE HALOGENS 319 

account. Hence, it may be omitted from the equation, only the 
oxygen, of which it is the source, appearing: 

I 2 + H 2 + 50 -» 2HI0 3 . 

In both these actions the initial substances (including the excess of 
nitric acid) and the products, with the exception of the iodic acid 
itself, are all volatile. When the solution is concentrated by evap- 
oration, therefore, only the iodic acid crystallizes. It is a white 
solid, perfectly stable at ordinary temperatures, and can be kept 
indefinitely. At 170° it begins to give off water vapor (2HI0 3 <=* 
H 2 + I 2 5 ), leaving iodic anhydride. The latter is a white 
crystalline powder which may be raised to 300° before it, in turn, 
breaks up, giving iodine and oxygen. 

Chemical Relations. — The compounds of the halogens with 
metals and with hydrogen diminish in stability, with ascending 
atomic weight of the halogen, in the order: F (19), CI (35.5), Br (80), 
I (127). Each halogen will displace those following it from this 
kind of combination. In the case of the oxygen compounds, the 
order of stability is just the reverse, those of iodine, for example, 
being the only ones which are reasonably stable. 

Amongst the oxygen acids of any one halogen, those containing 
most oxygen are most stable. The salts are in all cases more stable 
by far than the corresponding acids. 

The halogens when combined with metals and hydrogen are 
univalent (HI, KC1, etc.). It is clear, however, that, when united 
with oxygen, their valence is higher. The maximum is shown in 
perchloric anhydride (C1 2 7 ), where chlorine appears to be septi- 
valent. 

The formulae of the acids might be written so as to retain the 
univalence : 

H-Cl, H-O-Cl, H-O-O-Cl, H-O-O-O-Cl, 
H-O-O-O-O-Cl. 

But compounds in which we are compelled to believe that two oxy- 
gen units are united are usually unstable (e.g., hydrogen peroxide, 
H— O — O — H), and we should expect the instability would be 
greater with three and with four units of oxygen in combination. 
Here, however, the reverse state of affairs must be taken account 



320 COLLEGE CHEMISTRY 

of in our formulae, for HC10 4 is the most stable of the chlorine set. 
This reasoning, together with the septivalence in C1 2 7 , leads us 
to assume the valence seven in perchloric acid (see Periodic system) 
The structural formulae (cf. p. 292) of some of these substances are 
therefore written as follows: 



II II 

H-Cl, H-O-Cl, H-0-Cl = 0, Na-0-I = 0. 

II II 

.0 o 

Oxidation and Reduction 

Oxidation by Oxygen. — The simplest oxidations are the 
cases where a metal or non-metal unites with oxygen: 
2Cu + 2 -> 2CuO, S + 2 -> S0 2 . 
Union of a compound with additional oxygen is oxidation also. 
2S0 2 + 2 -> 2S0 3 , 3KC10 -> 2KC1 + KC10 3 . 

The removal of hydrogen from hydrogen chloride (preparation of 
chlorine, p. 156), is also defined as oxidation. 
2 + 4HC1 -> 2H 2 + 2C1 2 . 
2KMn0 4 + 16HC1 -> 8H 2 + 2KC1 + 2MnCl 2 + 5C1 2 . 

Every oxidation is accompanied by reduction of the oxidizing agent. 
Thus, in the second last equation, the free oxygen is reduced to 
water. Again, in the third last equation, 2KC10 is reduced to 
2KC1, while 1KC10 becomes KC10 3 by oxidation. 

In the laboratory, we frequently discover that an oxidation has 
occurred by noticing the presence of a product of reduction. 
Thus, when we heat carbon with sulphuric acid: 2H 2 S04 + C — > 
C0 2 + 2H 2 + 2S0 2 , we do not notice the product of oxidation, 
C0 2 , because it is odorless and colorless, but we perceive at once 
the odor of the sulphur dioxide, and realize that the sulphuric acid 
must have oxidized some substance, or this gas would not have 
been formed at the temperature employed. 

Note that the removal of the elements of water is neither oxida- 
tion nor reduction, for equivalent amounts of both oxygen and 
hydrogen are removed: 

2HC10 -> H 2 + C1 2 0, H 2 C0 3 -> H 2 + C0 2 . 



OXIDATION AND REDUCTION 321 

In the cases discussed above, oxidation consists always in adding 
oxygen or removing hydrogen. 

Oxidation by Other Negative Elements. — Oxygen is only 
one of the class of elements called non-metallic or negative ele- 
ments, so we cannot logically restrict the term "oxidation" to 
actions involving oxygen. Thus, forming a chloride, or increasing 
the proportion of chlorine in a compound is oxidation: 
Cu + Cl 2 -> CuCl 2 , 2FeCl 2 + Cl 2 -> 2FeCl 3 . 

In every compound, one of the elements is relatively positive 
and the other relatively negative. Thus, copper is positive and 
chlorine negative. In carbon dioxide C0 2 , carbon is (relatively) 
positive and oxygen negative, and in calcium carbide, CaC 2 , cal- 
cium is positive and carbon (relatively) negative. 

Thus, oxidation is introducing, or increasing the proportion of the 
negative element, or removing, or reducing the proportion of the 
positive element. Reduction is the converse. 

Oxidation and Valence. — Combining a metal with oxygen 
or sulphur raises the active valence of the metal from zero to some 
finite value: 2Cu° + 2 ° — » 2Cu n O n . Metallic copper has no 
valence in use. In CuO or CuCl 2 it has gained the valence II. 
The copper has been oxidized. Similarly, changing FeCl 2 into 
FeCl 3 increases the active valence of the iron from II to III (oxida- 
tion). Conversely, changing 2HC1 to Cl 2 decreases the active 
valence of chlorine from I to zero (oxidation). In the same 
equation (p. 320), KMn in KMnG 4 must have a total valence of 
VIII, but in the products KC1 + MnCl 2 the total valence has 
decreased to III (reduction). 

Again, in displacement, e.g., Zn + 2HC1 — > ZnCl 2 + H 2 , the 
zinc is oxidized because the active valence goes from zero to II, 
and the hydrogen is reduced. 

Hence, oxidation consists in increasing the active valence of a 
positive element or decreasing that of a negative element. Reduc- 
tion is the converse. 

This way of stating the rule makes it clear why removing the 
elements of water is neither oxidation nor reduction. We are 
removing both a positive and a negative element, and are removing 
them in equi-valent amounts, 2H 1 + O u . 



322 COLLEGE CHEMISTRY 

Oxidation and Ionization. — If, in the last illustration, we 
write the equation ionically: Zn + 2H+— >Zn ++ + H 2 , we dis- 
cover that, logically, we must consider the change from metallic 
zinc to zinc-ion to be in itself oxidation. This is the case whether 
the zinc-ion later combines with a negative ion to form a molecule 
or not. Mere union or disunion of ions is neither oxidation nor 
reduction. Conversely, the discharge of the 2H+ giving H 2 is 
reduction. 

Thus, ionization of an elementary substance to form a positive 
ion is oxidation, and ionization to form a negative ion is reduction, 
and conversely. 

Oxidation and Electrons. — Increasing the valence of an 
atom of a positive element (oxidation) consists in removing one or 
more electrons: Na° — e = Na + (p. 235). Increasing the valence 
of an atom of a negative element (reduction) means adding one or 
more electrons: Cl° + e — » Cl~. 

Hence, oxidation is removing electrons and reduction is adding 
electrons. 

Making Equations for Oxidations and Reductions. — The 

writing of equations for actions involving oxidation and reduction, 
where there are more than two substances on one side of the equation, 
is difficult, and a system or plan is of great value. The plan of 
partial equations (p. 194) is often helpful. There are three other 
systems which are in use. (1) When the action involves oxygen 
acids and their salts, the formulae can be rewritten so as to show 
the anhydride (see below). (2) The second, called the system of 
positive and negative valences, is more generally applicable (next 
section). (3) The third describes oxidation in terms of ions and 
positive electrical charges (p. 325). 

Making Equations: Using Positive and Negative Va- 
lences (p. 276). — 1. Each compound is composed of elements 
which are, relatively to one another, either positive or negative. 
Thus, in KMn0 4 , K and Mn are positive and is negative. In 
CS 2 , C is (relatively) positive and S negative (see p. 277). We 
say, then, that C has a positive valence of four (+4) and S has a 
negative valence of two ( — 2), just as it has in H 2 S. 



OXIDATION AND REDUCTION 323 

2. In each compound, the algebraic sum of the positive and negative 
valences must be zero. Thus, in CS 2 the sum is +4 — 2x2 = 
(C#S 2 = ). This is simply the rule of equi-valence (p. 62), with the 
addition of the idea of relative positiveness and negativeness. 

This enables us to determine the valence of each element in a 
compound like KMn0 4 . K + is always univalent and positive. = , 
in inorganic compounds, is always bivalent and negative. The 
valence of Mn has different values: Mn n Cl 2 , Mn 2 ra 3 , Mn IT 2 , 
Mn 2 vn 07, etc. By the rule (sum of valences equals zero) we can 
tell the valence of Mn in this compound. The valence of 4 
(40=) is -8. That of K is +1. That of Mn must therefore be 
+7 (KMn+ vn 4 ).* Again, in HC10 3 , the valence of 3 is -6, 
that of H is +1, therefore that of CI must be +5. Still again, in 
K 2 Cr 2 7 , the valence of 7 is —14, that of K 2 is +2, that of Cr 2 
is therefore +12, and that of Cr necessarily +6 (K 2 Cr 2 +VI 7 ) . 

3. Since rule 2 applies to every compound used or produced in 
a chemical change, it follows that when in a reaction the valence of 
an element changes in value, that of one or more of the other elements 
must also change, so as to maintain the equality of + and — valences. 
Thus, if one element loses in valence, to the extent of +6, some 
other element (or elements) must lose — 6, or gain +6. The gain 
(or loss) of one element must cancel the gain (or loss) of some other 
element. 

4. The valence of a free element, that is, its active valence, is 
zero. A free element is also neutral — neither positive nor nega- 
tive — because it is not combined with any other element. 

Illustration of rules 3 and 4. Thus, in the action for preparing 
chlorine with manganese dioxide (p. 158) : 

Mn0 2 + 4HC1 -> MnCl 2 + 2H 2 + Cl 2 , 

4H (4H+) has the valence +4 on both sides. On the left side, 
4C1 (4C1~) has the valence —4: on the right, 2C1 has the valence 
—2, and Cl 2 has the valence 0. So far as chlorine is concerned, 
there is a change from —4 to —2, or a difference of —2. Again, 
on the right, Mn has the valence +2, while on the left side 
it has the valence +4, a difference of +2. The two differences, 
— 2 and +2, cancel one another. Stated otherwise, manganese 

* The reader should write this, and other formulae discussed below, so as to 
show the valences thus: K+M ir[-j-{- +0 1 = (c/. p. 276). 



324 COLLEGE CHEMISTRY 

lost +2 and chlorine lost — 2, so that the other + and — valences 
still in use remained equal in number, and equi-valence was 
preserved. 

Balancing an Equation. Suppose we wish to balance the equa- 
tion for the decomposition of chloric acid HC10 3 . We ascertain, 
in the laboratory, that the products are perchloric acid HC10 4 , 
chlorine dioxide C10 2 , and water. 

Skeleton: HC10 3 -» HC10 4 + C10 2 + H 2 0. 

Since H + and = do not change in valence, only CI has been 
affected. On the left side, the valences are 3 = — 6, H = +1, CI 
therefore = +5.* On the right side, in HC10 4 , the total valence 
of oxygen is —8 and of hydrogen +1. That of CI is therefore 
+7. In C10 2 , the valence of 2 is —4, and that of CI there- 
fore +4. Thus, CI changes, from +5, partly to +7 and partly to 
+ 4. To achieve this, arithmetically, we require 3C1 on the left 
(= 3 X +5 = +15), giving CI = +7 and 2C1 = 2 X +4 = +8, 
or a total of + 15 on the right. Thus, we require 3HCIO3: 

Balanced: 3HC10 3 = HC10 4 + 2C10 2 + H 2 0. 

Balancing Another Equation. In the reaction for preparing 
chlorine (p. 157), the skeleton is: 

Skeleton: KMn0 4 + HC1 -» H 2 + KC1 + MnCl 2 + CI. 

Here, in KMn0 4 , the'valence of Mn is +7. In MnCl 2 it is +2, 
a loss of +5. The chlorine also changes its valence from — 1 to 0, 
a loss of —1. Evidently, so that the changes may cancel out, for 
every Mn losing +5, 5C1 must lose 5 X —1 and be liberated: 

Incomplete: KMn0 4 + HC1 -> H 2 + KC1 + MnCl 2 + 5C1. 

Since there is now, altogether, 8C1 on the right, 8HC1 will be re- 
quired on the left. The 8H will give 4H 2 0: 

Balanced: KMn0 4 + 8HC1 -> 4H 2 + KC1 + MnCl 2 + 5C1. 
Molecular: 2KMn0 4 + 16HC1 ^8H 2 + 2KC1 + 2MnCl 2 + 5C1 2 . 

For another method of balancing this equation, see p. 157. 

* Write these (and other formulae) thus: H+ Crft+Or (c/. p, 276). 



OXIDATION AND REDUCTION 325 

Making Equations, by Using the Anhydrides. — To balance 
the equation for the decomposition of chloric acid, we first write 
the skeleton equation: 

Skeleton: HC10 3 -> HC10 4 + C10 2 + H 2 0. 

Then we divide the acids into water and the anhydrides (p. 316). 
Analyzed: H 2 0,C1 2 5 -> H 2 0,C1 2 7 + C10 2 + H 2 0. 

We now perceive that, disregarding the water, some Cl 2 05 must 
lose oxygen to give 2C10 2 + O, and that some C1 2 5 must gain 20, 
becoming C1 2 7 . To furnish the 20, clearly 2C1 2 5 is required, 
giving 4C10 2 + 20, and a third C1 2 5 gains this 20. Thus, alto- 
gether 3C1 2 5 will be required: 

Balanced: 3H 2 0,C1 2 5 -> H 2 0,C1 2 7 + 4C10 2 + 2H 2 
or 6HCIO3 -> 2HC10 4 + 4C10 2 + 2H 2 0. 

This equation is then divided by two throughout. 

Making Equations by Oxidation of Ions, Using Positive 
Electrical Charges. — All oxidation reactions involving ionogens 
can be written in terms of ions. Thus, the oxidation of hydro- 
chloric acid by potassium permanganate can be so written. The 
potassium-ion clearly is not affected, and may be omitted. The 
ions concerned are: 

Mn0 4 ~ + H+ + CT -> H 2 + Mn++ + Cl°. 

Cl° with no charge stands for free chlorine. Now we can divide 
the action into (1) the behavior of the oxidizing agent, which 
is general, and will be used wherever the same oxidizing agent is 
used; (2) the fate of the substance being oxidized, which again is 
general, because other oxidizing agents will change it in the same 
way. 

MnOr + 8H+ -> 4H 2 + Mn++ + 5© . (1) 

In words, each permanganate ion, with a free acid present (oxi- 
dizing mixture), will give water, manganous-ion, and a balance 
of five unit positive charges. 

50 + 5CT -> 5CR (2) 

10© + 50 2 - -* 5H 2 + 50° 2 . (2 1 ) 

10© -f 5S0 3 = + 5H 2 -* 5S0 4 = + 10H+ (2 11 ) 



326 COLLEGE CHEMISTRY 

These three equations represent the oxidation of (2) hydrochloric 
acid, or (2 1 ) hydrogen peroxide, giving free oxygen, or (2 11 ) sul- 
phurous acid, with water furnishing the oxygen, and leaving the 
solution strongly acid (= 5H 2 S0 4 ). Note that the sums of the + 
and — charges on opposite sides of each equation are equal, 
To obtain the final ionic equation, add (1) and (2) : 

Mn0 4 " + 8H+ -> 4H 2 + Mn++ + 50 . (1) 

50 +5C1-->5C1°. (2) 

Mn0 4 ~ + 8H+ -f 5CP -> 4H 2 + Mn++ + 5CR 

Before adding (1) and (2 1 ) and (1) and (2 11 ), the first equation (1) 
must be doubled throughout, so that the 10© may cancel out. 

Exercises. — 1. Assign to its proper class (pp. 166, 258) each 
of the actions mentioned in this chapter. 

2. Knowing that potassium fluosilicate K 2 SiF 6 is insoluble, 
how should you make chloric acid (p. 314)? 

3. Make the equation for the interaction of chlorine with 
calcium hydroxide in hot water (p. 313). How should you make 
zinc chlorate from zinc hydroxide Zn(OH) 2 ? 

4. How should you make pure potassium hypochlorite from 
hypochlorous acid (p. 254)? 

5. Explain, in terms of ionic equilibrium, why dilute hypo- 
chlorous acid can be obtained by adding one-half of an equivalent 
of an active acid (p. 309) to bleaching powder, and distilling the 
mixture. 

6. On what circumstances would the possibility of making 
barium chlorate by action of chlorine on barium hydroxide depend 
(p. 313)? Could pure barium chlorate be obtained easily by this 
means (see Table of Solubilities)? 

7. Make the equations for: (a) the preparation of potassium 
bromate; (6) pure aqueous bromic acid; (c) the interaction of 
iodine with aqueous potassium hydroxide in the cold, and (d) 
when heated. 

8. Make the equations for the interactions of chlorine dioxide 
with water, and with aqueous potassium hydroxide. 

9. Find the formulae of the anhydrides of the following acids: 

HP0 3 , H 2 Se0 4 , H 3 As0 3 , H 3 As0 4 , H 6 S0 6 . 



OXIDATION AND REDUCTION 327 

10. Find the formulae of the anhydrides of the acids from the 
following formulae of salts: 

Na^SiOs, Na2HP0 4 , NaH 2 P0 3 , Na^HsIOe. 

11. Classify the following changes as oxidations or reductions, 
(a) H 2 Cr 2 7 -> H 2 Cr0 4 + Cr0 3 ; (6) HMn0 4 -> Mn0 2 ; (c)I ->I _ ; 
{d) 2H 2 2 -> 2H 2 0+0 2 . 

12. Using positive and negative valences, determine whether 
each of the following formulae is correct or incorrect: Ca(Mn0 4 ) 2 , 
A1(C10 4 ) 3 , Na*HI0 5 . 

13. Apply each of the three methods (pp. 322, 325) of writing 
equations to the four following reactions: (a) chlorine-water on 
bromine; (6) chlorine-water on hydrogen sulphide, giving free 
sulphur; (c) potassium permanganate and free acid on hydrogen 
sulphide, giving free sulphur; (d) potassium dichromate and free 
acid (p. 224) on hydrogen sulphide, giving the chromic salt of the 
acid (Cr m ) and free sulphur. 



CHAPTER XXIV 



THE ATMOSPHERE. THE HELIUM FAMILY 

The pressure which is exerted by the air upon each squaie centi- 
meter of the earth's surface is 1033.6 g., or a little over one kilo- 
gram. This is nearly fifteen pounds to the square inch. 

There are three classes of components in the air. Those of the 
first class, oxygen, nitrogen, and the inert gases of the helium 
family, are present in almost constant quantities. Those of the 
second class are very variable in quantity, and include carbon 
dioxide, water vapor, and dust. Those of the third class, such as 
the sulphur dioxide in city air, are accidental. 

Components which are Constant in Amount. — In the 

determination of the oxygen in air, phosphorus enclosed in iron 
gauze (Fig. 90), may be used. The oxygen com- 
bines to form several oxygen acids of phosphorus. 
The volume of gas is read off before the phosphorus is 
introduced, and after it has been withdrawn. 

In the air taken from mines, from mountain tops, 
from the surface of the sea, and from inland regions, 
the percentages of oxygen by volume are found to be 
very constant (20.9 to 21.0). 

When air, from which the oxygen has been re- 
moved by phosphorus, or by passage over heated 
copper or iron, is led slowly through a heated tube 
containing magnesium, the nitrogen unites with the 
metal to form the solid magnesium nitride Mg 3 N 2 , 
and only about 10 c.c. out of every liter remains 
uncombined. This residuum is argon, mixed with 

0.15 per cent of its volume of other gases belonging to the helium 

family. 

The Carbon Dioxide. — Pure country air contains about 3 
parts in 10,000 of carbon dioxide C0 2 . In city air there are from 6 

328 




Fig. 90. 



THE ATMOSPHERE 329 

to 7 parts in the same volume, while in the air of audience-rooms 
the proportion may rise as high as 50 parts. 

The sources of the carbon dioxide in the air are numerous. It 
comes from the decay of vegetable and animal matter, in which, 
chiefly through the influence of minute vegetable organisms, the 
carbon is oxidized to carbon dioxide. It is formed also by the com- 
bustion of coal and wood, but the thirteen hundred million tons 
of coal burned annually, giving three times that weight of carbon 
dioxide, would add only one-six hundredth to the total present in 
the air. It is exhaled by animals, being produced in the body by 
oxidation of the carbon in the food which they eat. It also issues 
from the earth, in volcanic as well as in other neighborhoods. The 
proportion of this gas in the air would naturally increase continu- 
ously, though slowly, as the result of these processes, were it not 
that it is removed just as continuously by the action of growing 
plants (see p. 387), which use it as food. It may be added, also, 
that carbon dioxide, being a soluble gas, is contained in sea water 
(dissolved and as Ca(HC0 3 )2), and the total amount in the ocean 
is much greater than that in the air. The removal by plants and 
by sea water thus keeps the proportion in the air fairly constant. 

The presence of carbon dioxide in the breath may be shown 
very quickly by blowing through a tube into calcium hydroxide 
solution (limewater). Calcium carbonate CaC0 3 is precipitated. 
We draw about 500 c.c. of air into our lungs at each breath, or 
half a cubic meter per hour. In the lungs, some oxygen is re- 
moved, the percentage by volume falling from 21 to 16, and we 
add some carbon dioxide, the proportion increasing from 0.03 in 
country air to about 4 per cent. A candle flame is extinguished 
by exhaled air, because the maintenance of such a flame requires 
at least 18.5 per cent of oxygen. But air will sustain life until the 
proportion has fallen to about 10 per cent. 

To determine the proportion of carbon dioxide, a measured volume 
of air is bubbled slowly through a measured volume of a solution 
of barium hydroxide of known concentration. Barium carbonate is 
precipitated: Ba(OH) 2 + CO2 — > BaC0 3 J, + H 2 0, and the quantity 
of barium hydroxide remaining is determined by titration (p. 256). 

The Water Vapor. — The proportion of water vapor is con- 
stantly changing. When the air becomes cool, as it does most often 



330 COLLEGE CHEMISTRY 

in the upper layers, the vapor condenses to droplets, forming fogs 
and clouds. When the condensation continues, the drops be- 
come larger and fall as rain. On the other hand, when the weather 
is warm, water from the soil, and from rivers, lakes, and oceans, 
passes into vapor and the amount in the air increases. 

Humidity. — The moisture in the air is usually defined in 
terms of the relative humidity, the standard being the quantity 
required to saturate the air. The open air is never actually 
saturated, but, when a portion is confined in a vessel over water, 
it soon becomes so. The humidity is then 100 per cent. If the 
partial pressure of water vapor present is only half as great as the 
vapor pressure of water at the same temperature, the humidity is 
50 per cent. The average humidity is roughly about 66 per cent. 

At 18° (64.4° F.), the vapor pressure of water is 15.4 mm. Thus 
air saturated with moisture at 18° (100 per cent humidity) would 
contain 15.4/760, or about 2 per cent by volume of water vapor. 
If this air were cooled to 0° (32° F.), a temperature at which the 
vapor pressure of water is only 4.6 mm., the air could retain only 
4.6/760, or 0.6 per cent, of moisture. The difference, amounting 
to 10.4 g. (10.4 c.c.) of water per cubic meter, would condense as 
fog or rain. 

The proportion of water in a given volume of air may be meas- 
ured most accurately by permitting the air to stream slowly 
through tubes filled with calcium chloride or phosphoric anhydride. 
The increase in weight of the charged tubes represents the quantity 
of moisture abstracted from the sample. It may also be ascer- 
tained by noting the temperature to which air has to be cooled 
before it becomes saturated and deposits dew (dew-point). For 
example, if air at 18° has to be cooled to 11° before it deposits dew, 
it contains water vapor at a pressure of 9.8 mm. (Appendix IV). 
If saturated at 18°, it would have contained water vapor at a 
partial pressure of 15.4 mm. The relative humidity was, there- 
fore, 9.8/15.4, or 63.6 per cent. 

Ventilation. — On a moist day, we speak of the atmosphere 
as "heavy" or "oppressive." The barometer, however, is lower 
on such days, and the pressure below the average. Moist air 
must be lighter than dry air, because in moist air molecules of 



THE ATMOSPHERE 331 

relative weight 18 (H 2 0) have been substituted for an equal num- 
ber of molecules of oxygen and nitrogen with the relative weights 
32 and 28. The discomfort is due to a different cause. 

The oxidation of digested food carried by the blood is accom- 
panied by liberation of heat, yet our bodies must remain at 98.6° F. 
(37° C). A rise of a few tenths of a degree produces discomfort. 
A little of the heat is lost by radiation from the surface of the body 
but the real adjustment is secured by evaporation of water through 
the skin. The vaporization of 1 g. of water (at 100°) removes heat 
amounting to 540 calories (603 cal. at 37° C). Evaporation of a 
single ounce (28-| g.) of water will therefore lower the temperature 
of 96.5 kilograms (168 lbs.) of water (or flesh, which is largely 
water) by more than two-tenths of a degree C. (nearly 0.4° F.). 

The " oppressive" feeling, then, is due to the fact that the air 
is too nearly saturated, evaporation is being hindered (p. 90), 
and heat is accumulating. Hence, the relative humidity is the 
measure of the goodness or badness of the air of a room. 

In winter, cold and therefore relatively dry air is brought into 
the house and heated. This makes the relative humidity very 
low, evaporation proceeds too fast, and discomfort follows. In 
summer, however, the outside air is often already nearly satu- 
rated at the temperature of the room. Unless there is a rapid 
change of air by ventilation, the moisture from the bodies of those 
in the room increases the humidity, and discomfort arises from a 
cause opposite to the one which produced it in winter. 

It should be noted, also, that even though the air is in constant 
motion, the layer of air next our skin (even the exposed parts) is 
hindered from moving by friction. There is a stationary layer 
close to the surface, which quickly reaches the temperature of the 
body and becomes saturated at that temperature. The water 
molecules can leave this layer, and make room for others, only by 
diffusion, which is a deliberate rather than a speedy process. 
Now, an electric fan, although it brings no fresh, dryer air into the 
room, nevertheless stirs the air and blows away the moist, saturated 
layer next the skin. It, at least, makes this layer much thinner, 
and reduces greatly the distance the water molecules have to go 
by mere diffusion.* 

* The same conception applies to dissolving a salt. A stationary layer of 
saturated solution is formed on the surface, and the molecules of the salt can 



332 COLLEGE CHEMISTRY 

Formerly, the accumulation of carbon dioxide from the breath 
was blamed for the unhealthiness of unventilated rooms. The 
proportion found in such rooms, however, is almost never sufficient 
to do any harm. Then, it was imagined that traces of highly 
poisonous compounds were exhaled by the body. No one, how- 
ever, has yet been able to prove that such poisons exist. 

The aims of ventilation are, therefore, to supply fresh outside 
air, to keep it in motion, and to maintain a humidity that is 
neither too low nor too high. 

Dust in the Air. — A beam of sunlight, crossing a dark room, 
can be seen by the light reflected from the particles of dust in the 
air. Some of the particles are inorganic, and consist of clay, 
limestone, and soot from ill-burned fuel. The organic dust may 
be divided into two kinds. The part which is dead includes coal 
dust, refuse from the streets, minute shreds of cotton, linen, hay, 
etc. The living dust consists of pollen grains, spores of fungi 
and molds, bacteria, and similar microscopic organisms. The 
presence of microscopic germs in the air is shown by the fact that 
when food has been exposed to the air, even for a few minutes, 
putrefaction very soon sets in. Some germs also produce disease 
when they land on a place where the skin has been damaged by a 
cut or a burn. After infection, antiseptic treatment, e.g., with 
hydrogen peroxide, destroys the organisms. But protection, e.g., 
with petrolatum (p. 391), until a new skin has formed, is better. 

It is worth noting that natural soil contains about 100,000 
micro-organisms per c.c, good, unfiltered river water from 6000 to 
20,000 per c.c, and pure air only 4 or 5 per liter. 

If dust were absent from the air, there would be no clouds or 
rain. Aitken has shown that the water vapor will not condense to 
fog in air that has been freed from dust by filtration. When moist 
air is cooled, the dust particles act as nuclei, round which the 
liquid grows at the expense of the vapor. In the absence of dust, 
the cooling would produce supersaturation, which would be slowly 

escape, and make room for more, only by diffusion. In liquids, this is a very 
slow process. By shaking the solid and liquid, however, the stationary layer 
is partly washed away. It is made thinner, so that the distance the mole- 
cules have to travel by diffusion is greatly reduced, and the whole operation 
is hastened. 



THE ATMOSPHERE 



333 



relieved by condensation on the surfaces of houses, plants, animals, 
and land. Thus, in a dustless atmosphere an awning or umbrella 
would afford no shelter. 

The formation of fog in ordinary air, and its absence in filtered 
air — e.g., air drawn through a wide tube packed with 20-30 inches 
of cotton — is easily shown in a 
darkened room (Fig. 91). The 
flask contains some water to satu- 
rate the air. When suction is 
applied, by the mouth, to the tube 
S, the saturated air in the flask 
expands and is cooled.* With ordi- 
nary air, a fog, brilliantly illumi- 
nated by the beam of light, is 
instantly produced. Filtered air 
(dustless) gives no fog. On the 
other hand, a whiff of smoke from 
smoldering paper, when admitted to the flask, causes a fog (after 
cooling) of extraordinary denseness. 




Fig. 91. 



Composition of Air. — Air, when freed from carbon dioxide 
and water, contains by volume 78.06 per cent of nitrogen, 21.00 
per cent of oxygen, and 0.94 per cent of argon. When only the 
water is removed, the carbon dioxide averages about 0.03 per cent 
of the whole. 

To use an illustration of Graham's, if we imagined the air to be 
divided by magic into layers, all at one atmosphere pressure, and 
with the heavier components below, we should have: On the earth, 
five inches of water; above that, thirteen feet of carbon dioxide; 
above that, ninety yards of argon; above that, one mile of oxygen; 
and on the top four miles of nitrogen. 



Air a Mixture. — The experiments, in which the oxygen was 
removed from the air and the nitrogen remained, do not prove 
that the original constituents were present simply in mechanical 
mixture. They might have been combined, and the combustion 
of phosphorus, for example, might have represented the removal of 
oxygen from combination with nitrogen and its appropriation by 
* Compression with a bicycle pump heats air, and expansion cools it. 



334 



COLLEGE CHEMISTKY 



the phosphorus. It may be well, therefore, to point out some 
reasons which lead us to regard the air as a mixture: 

1. Each of the substances in air has precisely the same properties 
which it exhibits when free, separate, and pure. This is char- 
acteristic of a mixture. Thus, the density of air is precisely that 
which we find by calculation from the known proportions and 
several densities of the components. Again, the solubility of each 
gas is observed to be the same as if the same amount of it were 
present, alone, in the same volume. 

2. When liquefied air is allowed to evaporate in a suitable 
apparatus, the nitrogen, being more volatile, can be separated 
completely from the oxygen. When the oxygen, in turn, is allowed 
to evaporate, the carbon dioxide and water remain as solids, 

frozen at this low temperature. 

3. Finally, the proportions by weight 
cannot be represented by a chemical 
formula, because they are not exact 
multiples of the atomic weights by 
integral numbers. This is a sure proof 
that it is not a chemical aggregate. 

Liquefaction of Gases. — The 

earliest experiments of this kind were 
made by Northmore (1805), who lique- 
fied chlorine, hydrogen chloride, and 
sulphur dioxide. In 1823 chlorine was 
again liquefied by Faraday. During 
the following years he reduced sulphur 
dioxide, hydrogen sulphide, carbon di- 
oxide, nitrous oxide, cyanogen, and 
ammonia to the liquid condition. He 
failed, however with oxygen, hydrogen, 
and nitrogen. In 1883 Wroblevski and 
Olszevski prepared visible amounts of 
liquid oxygen. About the same time Dewar devised means of manu- 
facturing large quantities of liquid air and oxygen. The most suc- 
cessful apparatus for use on a small scale is that devised by Hampson. 
In Hampson's apparatus (Fig. 92), two concentric copper pipes, 
about 130 meters in length, are coiled closely in a cylindrical form, 




Fig. 92. 




THE HELIUM FAMILY 335 

with non-conducting covering to prevent access of heat. Air at 
130-150 atmospheres pressure is forced through the inner pipe. 
When it reaches the extremity of this pipe, it suddenly escapes into 
a closed vessel. This expansion lowers its temperature. The air 
can now escape only by traveling back through the outer pipe to 
the final exit near the top. In doing so, it cools the highly com- 
pressed air in the inner pipe. This cooler air, on reaching the 
closed vessel, expands and becomes colder than ever, and in 
passing backwards lowers the temperature of the air in the inner 
pipe still further. Finally, the air in this pipe liquefies, and drops 
of liquid air are expelled into the closed vessel. They are allowed 
to run out through a valve, from time to time, as they accumulate. 

The cooling on expansion depends upon the imper- 
fection (p. 78) of the gas, and is due to the work 
done in overcoming the tendency to cohesion of its 
molecules. Liquid air can be kept in Dewar flasks 
(Fig. 93). The space between the inner and outer 
flasks is evacuated, so that there is no gas to carry 
heat to the liquid air. The inner surface of the outer 
flask is often silvered, so that radiant heat, from surrounding 
bodies, may be reflected and not absorbed. 

Liquid Air. — Liquid air varies in composition, as the nitrogen 
(b.-p. —194°) is less condensible than the oxygen (b.-p. —181.4°). 
It boils at about — 190°, and contains about 54 per cent of oxygen 
by weight, while air contains 23.2 per cent. By allowing evapora- 
tion to go on, a liquid containing 75 to 95 per cent of oxygen is 
easily obtained (cf. p. 26). The gas secured by the evaporation 
of the residue is pumped into cylinders and sold as compressed 
oxygen. It contains about 3 per cent of argon, and is a con- 
venient source of this element. Cartridges made of granular 
charcoal and cotton waste, when saturated with liquid air, have 
been used as an explosive in mining. 

The Helium Family 

Argon A. — Cavendish (1785) sought for other gases in air by 
adding more oxygen, passing an electric discharge to cause this 
gas to combine with the nitrogen, and absorbing the product 
(NO a ) in potassium hydroxide solution. He found that only 



336 COLLEGE CHEMISTRY 

about 0.8 per cent of inactive gas remained. Since the quantity 
was so small, and the spectroscope, by which the gas even in small 
amounts would have been recognized to be new, was not invented 
until much later, he did not pursue the subject. 

A century later, Lord Rayleigh observed that, while specimens 
of oxygen and other gases made purposely from various sources 
always had the same density, nitrogen was an exception. One 
liter of nitrogen made from air, and supposed to be pure, weighed 
1.2572 g. When the gas was manufactured by decomposition of 
five different compounds, such as urea and certain oxides of nitro- 
gen, the mean weight of a liter of this nitrogen was only 1.2505 g. 
The difference, amounting to nearly 7 mgm., was very much greater 
than the experimental error. The suspicion naturally arose that 
some heavier gas was present in natural nitrogen. Soon after 
(1894), Rayleigh repeated Cavendish's experiment, and obtained 
argon. Working in cooperation with him, Professor, now Sir 
William Ramsay, obtained the same gas by removal of the greatly 
preponderating nitrogen by means of magnesium (p. 328). The 
new gas had a molecular weight of about 40, and was therefore 
more than one-third heavier than nitrogen. 

The exact density of argon is 39.88. When liquefied it boils at 
- 186.9°, and the colorless solid melts at - 189.5°. The solubility 
of the gas in water (4 volumes in 100) is two and one-half times 
that of nitrogen. It has not been found to enter into any sort of 
chemical combination, and was named argon on this account (Gk., 
inactive). The physical properties show that the molecules of the 
gas, like those of mercury (p. Ill), are monatomic. 

Helium He, — In 1868 Lockyer first detected an orange line 
in the spectrum of the sun's prominences which was not given by 
any terrestrial substance then known. The line was so con- 
spicuous that it was attributed to the presence of a new chemical 
element, which was named helium (Gk., the sun). Ramsay, in 
searching for sources of argon, examined a gas which Hillebrand 
had obtained from uraninite, an ore of uranium. He was sur- 
prised to find (1895) that it contained a large proportion of a very 
light gas, the spectrum of which was identical with that of solar 
helium. The same gas is found in small amount in the atmosphere. 
Helium does not exhibit any tendency to enter into combination. 



THE HELIUM FAMILY 337 

It is monatomic and its density shows that its molecular weight is 
4. When liquefied by Onnes, it boiled at -268.5° (4.5° Abs.). 

Neon Ne, Krypton Kr, and Xenon Xe. — By liquefying 
atmospheric argon, using liquid air to cool it, and distilling the 
liquid, Ramsay (1898) found that, besides argon, it contained 
helium, along with three new gases. These together constituted 
one-six hundredth part of the whole. The gases were named neon 
(Gk., new), krypton (Gk., hidden), and xenon (Gk., stranger). 
These gases are all entirely inactive chemically, and are all mona- 
tomic. Their molecular weights are: Neon, 20.2; krypton, 82.9; 
xenon, 130.2. Niton Nt (radium emanation, q.v.), molecular 
weight 222.4, also belongs to this family. 

Exercises. — 1. A sample of moist air, confined over water at 
15° and 760 mm., occupies 15 c.c. It is mixed with 20 c.c. of hydro- 
gen, and the mixture is exploded, and suffers a contraction of 9.5 
c.c. What would be the volume of the oxygen it contained if 
measured dry at 0° and 760 mm.? 

2. Calculate, from the data on p. 333 and the densities, the 
percentage b} r weight of the three principal components of air. 

3. Of the proofs that air is a mixture (p. 333), which show that 
no part of the components is combined, and which that the com- 
ponents are not wholly combined? 

4. What is the relation between heavier clothing and the 
stationary layer of air next the skin? 

5. From the data given on p. 330, calculate the weight of water 
vapor in 1 cubic meter of air saturated at 18° and at 0°, respectively. 



CHAPTER XXV 

NITROGEN AND AMMONIA 

Nitrogen was recognized to be a distinct substance by Ruther- 
ford (1772), Professor of Botany in the University of Edinburgh, 
who named it mephitic air. Scheele showed that it was present 
in the atmosphere. Lavoisier recognized it to be an element, and 
named it azote (Gk., without life) because it did not support life. 
The English name records the fact that it is an important con- 
stituent of niter KN0 3 . 

The Chemical Relations of the Element Nitrogen. — In 

compounds with hydrogen and the metals nitrogen is trivalent, 
while in those containing oxygen and other negative elements, it 
is frequently quinquivalent. It is a non-metal, for its oxides are 
acidic (p. 94). Many of the compounds of nitrogen are extremely 
active and interesting. Those of them which we have to discuss 
in inorganic chemistry are ammonia NH 3 and nitric acid HN0 3 , 
and several related substances. 

Occurrence. — Free nitrogen is present in the air. The 
nitrates of potassium and sodium are found in Bengal and Chile, 
respectively. Natural manures, such as guano, contain large 
quantities of nitrogen compounds, and owe their value as fertilizers 
to this fact. Nitrogen is a constituent of the proteins (about 15 
per cent nitrogen) of vegetable and animal matter. 

Preparation. — Nitrogen containing about one per cent of 
argon is obtained by burning phosphorus in air, or by passing air 
over heated copper: 2Cu + 2 — » 2CuO. For commercial pur- 
poses, it is obtained by evaporation of liquid air. 

Pure nitrogen is prepared by heating ammonium nitrite: 

NH4N0 2 ->2H 2 + N 2 . 
338 



NITROGEN AND AMMONIA 339 

In practice, since ammonium nitrite is unstable and cannot be 
kept as such, strong solutions of ammonium chloride and sodium 
nitrite are mixed, a double decomposition results in the formation 
of ammonium nitrite, NH4CI + NaN0 2 «=* NH4NO2 + NaCl, and 
this breaks up when heat is applied, giving nitrogen. 

We may also prepare nitrogen by the oxidation of ammonia 
NH 3 , passing the latter over heated cupric oxide (see p. 343), or 
by the reduction of nitric oxide NO, passing this gas over heated 
copper. 

Physical Properties, — Nitrogen is a colorless, tasteless, 
odorless gas, as we should expect from the fact that air possesses 
these properties. It forms a colorless liquid, boiling at —194°, 
and a white solid (m.-p. —214°). The solubility in water (1.6 
vols, in 100) is less than that of oxygen. The density of the gas 
shows the formula of free nitrogen to be N2. 

Chemical Properties. — Nitrogen unites with few elements 
directly. At ordinary temperatures it is almost absolutely in- 
different. When passed over heated lithium, calcium, magnesium, 
or boron, it forms nitrides, in which it is trivalent. These have 
the formulae Li 3 X, Ca 3 N 2 , Mg 3 N 2 , and BN, respectively. Thus, 
when magnesium is burned in the air, the white mass which is 
formed contains magnesium nitride, along with much of the 
oxide. When the ash is moistened with water in a covered vessel, 
ammonia can be smelt and can be detected with moist litmus 
paper. The nitride is hydrolyzed: 

Mg3N 2 + 6H 2 -» 3Mg(OH) 2 + 2NH 3 T. 

Nitrogen combines with difficulty with hydrogen to form am- 
monia NH 3 and with oxygen to form nitric oxide. The actions 
will be discussed under the compounds themselves. 

One case of direct union of nitrogen is of economic importance. 
The supply required by most plants is obtained from nitrogen 
compounds contained in fertilizers, or equivalent substances 
already present in the soil. With the leguminosce (peas, beans, 
clover, etc.), however, are found associated certain bacteria, 
which flourish in nodules upon their roots. These bacteria have 
the power of taking free nitrogen from the air, which penetrates 



340 COLLEGE CHEMISTRY 

the soil, and producing proteins. The nodules often contain 
over five per cent of combined nitrogen. The proteins, by the 
action of nitrifying bacteria, give nitric acid which, with bases in 
the soil, gives nitrates. These are soluble, and are absorbed 
through the roots, furnishing the nitrogen needed by plants to 
enable them to construct the proteins they require. 

Compounds of Nitrogen and Hydrogen. — The commonest 
and longest known of these substances is ammonia NH 3 , which 
was first described by Priestley (1774) and named " alkaline air." 
Curtius discovered hydrazine N2H4 in 1889, and hydrazoic acid 
HN 3 in 1890. Hydroxylamine HONH 2 , discovered by Lossen in 
1865, is similar to ammonia in chemical behavior. 

Ammonia NH 3 

Ammonia is of interest, commercially, because large amounts 
of liquefied ammonia are used in refrigeration, because much is 
employed in the manufacture of carbonate of soda, and because its 
compounds are used as fertilizers. 

Manufacture. — Ammonia is formed when proteins are heated 
in the absence of air. Thus, it was formerly obtained by the 
distillation of hoofs, hides, and horns, and the solution in water 
was called "spirit of hartshorn." Coal contains about 1 per cent 
of combined nitrogen, derived from the proteins of the original 
plants. Hence, when coal is distilled in the manufacture of coal 
gas or, on a much larger scale, for the making of coke, much am- 
monia can be secured by washing with water the gases which are 
given off. The solution is separated from the tar, lime is added 
to combine with acids, and the ammonia gas is driven out by 
heating and passed into sulphuric acid or hydrochloric acid. It 
gives ammonium sulphate or chloride (see below). 

In Germany 80 per cent (1910) of the coke is made in "by- 
product" coke ovens, in which the ammonia and other by-products 
are collected and utilized; in the United States 83 per cent of the 
coke is made in "beehive" ovens, in which the vapors are simply 
burned. Ammonium sulphate is a valuable fertilizer and in 
1911, in the United States, ammonia capable of yielding 400,000 



NITROGEN AND AMMONIA 341 

tons of ammonium sulphate worth 24 million dollars was burned 
by the cokemakers. 

The distillation of coal is the chief source of commercial am- 
monia. In Scotland, however, oil-bearing shale is distilled to 
obtain petroleum, and much ammonia, liberated at the same 
time, is collected. Formerly it was allowed to escape but, in the 
absence of a protective tariff, the competition of American and 
Russian petroleum compelled economy. Now, the profit on the 
ammonium sulphate pays the whole cost of mining and distilling 
the shale. 

Synthetic Ammonia, — The Badische Company is now 
manufacturing ammonia on a large scale, for the preparation of 
explosives, by the direct union of nitrogen and hydrogen. 

N 2 + 3H 2 *± 2NH 3 + 2 X 12,200 cal. 

No union occurs at low temperatures and, on the other hand, the 
action is reversible and exothermal, so that at 700° ammonia is 
decomposed almost completely (Van't Hoff's law, p. 188). It is 
necessary, therefore, to use a lower temperature and a contact 
agent — such as specially prepared iron — to hasten the action. 
Then, too, the reaction is accompanied by a diminution in vol- 
ume (4 vols. — > 2 vols.), and is therefore assisted by using the 
gases under a pressure of 185-200 atmospheres (Le Chatelier's 
law, p. 190). At 500°, with these conditions, about 8 per cent 
of the gases combine. The ammonia is dissolved out with water, 
and the uncombined gases are sent through the process again. 
The required hydrogen may be obtained by one of the com- 
mercial processes (p. 56), and the nitrogen from liquid air. 

Preparation in the Laboratory. — 1. A mixture of slaked 
lime and some salt of ammonium, such as ammonium chloride, 
either with or without water, is heated in a flask or retort pro- 
vided with a delivery tube: 

Ca(OH) 2 + 2NH4CI *=► CaCl 2 + 2NH40H <± 2NH 3 + 2H 2 0. 

The ammonium hydroxide, formed by the double decomposition, 
immediately decomposes. 



342 



COLLEGE CHEMISTRY 



2. Warming the aqueous solution gives a steady stream of the 
gas. Since the gas is very soluble in water, it is collected over 
mercury or in an inverted jar by downward displacement of air. 
In both methods of preparation, it is dried with quicklime (p. 475). 

Physical Properties. — Ammonia is a colorless gas with a 
pungent, characteristic odor familiar in smelling-salts. The 
G.M.V. of the gas weighs 17.26 g., so that the density is little 
more than half that of air (c/. p. 101). When liquefied it boils 
at —33° and the solid is white and crystalline (m.-p. —77°). 
One volume of water dissolves 1300 volumes of the gas at 0°, 
and 783 volumes at 16°. The 35 per cent solution, sold as "con- 
centrated ammonia," has a sp. gr. 0.881. The whole of the dis- 
solved gas may be removed by boiling (c/. p. 145). 

Liquefied ammonia is used in refrigeration. In evaporating at 
— 33° it absorbs 330 cal. per gram. Water alone has a greater 

heat of vaporization. The large amount 
of heat is, in both cases, required be- 
cause of the relatively large volume of 
the vapor (due to low molecular weight) 
and to the fact that both liquids are 
associated (p. 206), and the complex 
molecules (NH 3 ) 2 and (NH 3 )3 have to 
be decomposed. To freeze 1 gram of 
water at 0°, 79 cal. have to be re- 
moved. Thus 1 g. of liquid ammonia 
will convert 4 g. of water into ice. Fig. 
94 shows one arrangement diagram- 
matically. The ammonia gas, obtained 
from a cylinder of liquid ammonia, is 
driven by the pump F along the tube 
E and is liquefied in the tube coiled in 
the tank AB. Cold water circulating 
through AB removes the heat produced 
by the compression and liquefaction of the gas. The liquid 
ammonia is allowed to drip through the stopcock G into the lower 
coil, and there it evaporates. In doing so, it takes heat from a 
30 per cent solution of calcium chloride in water. This cooled 
brine leaves the tank at D, circulates through another tank, in 




Fig. 94. 



NITROGEN AND AMMONIA 343 

which water-filled ice molds are suspended, and returns to C. 
When used for cooling storage-rooms for meat, the brine circulates 
through pipes in the same way. The machine is constructed of 
iron, because copper and brass are corroded by ammonia. 

Chemical Properties. — Ammonia, as we have seen, is not 
stable, and decomposes almost completely at 700°. A discharge 
of sparks from an induction coil (temperature about of ^ 

2000°) has the same effect, so that a sample of the 
gas, confined over mercury in a closed tube (Fig. 95), 
may be shown to double in volume. Every two 
molecules give four: 

That, even at this temperature, the action, being 
reversible, is still incomplete, can be shown by 
introducing a few drops of dilute sulphuric acid. 
The trace of ammonia remaining combines with this 
acid, forming (XEL^SC^ in solution. If the dis- 
charge is continued, further traces of ammonia are 
formed and absorbed, until, finally, the whole gas 
disappears. 

Ammonia reduces many oxides, when the latter 
are heated and the gas is led over them: 

3CuO + 2NH 3 -» 3Cu + 3H 2 + N 2 . 

Ammonia burns in pure oxygen (not in air) to give 

steam and nitrogen. FlG - 95 - 

Chlorine and bromine (vapor) combine with the hydrogen and 
liberate nitrogen: 

2NH-, + 3C1 2 -> N 2 + 6HC1. 

When metals capable of uniting with nitrogen (p. 339) are 
heated in a stream of ammonia gas, hydrogen is displaced. Mag- 
nesium gives magnesium nitride: 

2NHa + 3Mg -> Mg3N 2 + 3H 2 . 

Sodium and potassium, however, give amides (compounds con- 
taining the group XH 2 ), such as sodamide NaNH 2 : 
2NH 3 + 2Na -+ 2NaNH 2 + H 2 . 



344 COLLEGE CHEMISTRY 

The most striking property of ammonia is that it combines with 
acids, giving ammonium salts: 

NH 3 (gas) + HC1 (gas) -> NI^Cl (solid) . 
2NH 3 (gas) + H 2 S0 4 (liq.) -> (NH4) 2 S0 4 (solid). 

It combines also with water at or below —79.3° to give ammonium 
hydroxide, a white solid: 

NH 3 + H 2 -> NH4OH. 

As the solid dissociates above —79.3°, a solution of the substance, 
which is contained in the aqueous solution of ammonia, is the 
only available form of ammonium hydroxide. In solution, it is a 
weak base. 

Ammonium oxide (NEL^O, a solid, can also be formed below 
-78.6°. 

Ammonium Compounds. — Since NH4 plays the part of a 
metallic element, entering into the composition of a base and of 
a series of salts, it is named ammonium. As this radical forms a 
univalent, positive ion NHf 1 " and gives a distinctly alkaline base, 
it is classed with the metallic elements of the alkalies (q.v.). 

Ammonium Hydroxide. — Although less completely ionized 
than potassium hydroxide, ammonium hydroxide affects litmus 
easily. In a normal solution, 0.4 per cent of the ammonia is in 
the form of ammonium-ion NH4+. When an acid is added to the 
solution, the equally small amount of hydroxide-ion which exists 
in it is removed and the various equilibria are displaced forward. 
The final result is the same as with any other base: 

NH 3 (gas) ^NH 3 (dslvd)+H 2 0<=>NH 4 OH±^NH4++OH-^ TT n 

HClt^Cl" + H+)-> n2U * 

Only a small proportion of the gas (one-third) is actually com- 
bined at any one time, the greater part being simply dissolved. 

The solution is sold as household ammonia, and is used, in 
washing and cleaning, to soften the water. 

Salts of Ammonium. — When strongly heated, all am- 
monium salts are decomposed and many, but not all, give am- 



NITROGEN AND AMMONIA 345 

monia and the acid. When the latter is volatile, the whole 
material of the salt is thus converted into gas. The acid and the 
ammonia reunite to form the solid salt when the vapor reaches a 
cool part of the tube (sublimation, p. 199) : 

NH4CI (solid) <=± NH4CI (gas) <± HC1 + NH 3 . 

The test for ammonium salts is to warm them, dry or in solution, 
with a base, when the odor of ammonia becomes noticeable. 

(NI 2KOh52k" + 2OT 1+ l^ 2NH * OH?±2H20 + 2NH 'f • 

When the solution is used, it is the tendency of the NH4 4 " and 
OH~~ to unite to form the slightly ionized, molecular hydroxide 
that sets the other equilibria in motion. 

In ammonium salts, the nitrogen is quinquivalent. 

Hydrazine N 2 H^. — By reduction of a compound of nitric oxide 
and potassium sulphite by means of sodium amalgam,* a solution 
of hydrazine hydrate is obtained: 

K 2 S0 3 ,2NO + 3H 2 -> N 2 H4,H 2 + K 2 S0 4 . 

When the hydrate is distilled with barium oxide, under reduced 
pressure, hydrazine is liberated: 

N 2 H4,H 2 + BaO -> NAT + Ba(OH) 2 . 

Hydrazine hydrate freezes at about —40° (b.-p. 118.5°). Its 
aqueous solution is alkaline, and salts are formed by neutraliza- 
tion. 

Hydrazoic Acid HN 3 . — When nitrous oxide (q.v.) is led over 
sodamide at 200°, water is liberated and sodium hydrazoate re- 
mains behind: 

NH 2 Na + N 2 -> NaN 3 + H 2 0. 

A dilute solution of the free acid is best obtained by distilling the 
lead salt with dilute sulphuric acid. 

The pure acid (b.-p. 37°) is violently explosive, resolving itself 
into nitrogen and hydrogen with liberation of much heat: 
2HN 3 , Aq -> H 2 + 3N 2 + Aq + 2 X 61,600 cal. 

* The sodium dissolved in the mercury interacts with the water, giving 
hydrogen (see Active state of hydrogen). 



346 COLLEGE CHEMISTRY 

Halogen Compounds of Nitrogen. — When ammonium 
chloride solution is treated with excess of chlorine, drops of an 
oily liquid, nitrogen trichloride, are formed: 3C1 2 + NH4CI — » 
NCI3 + 4HC1. It is extremely explosive, resolving itself into its 
constituents with liberation of much heat. 

When a solution of iodine in potassium iodide solution (p. 200) 
is added to aqueous ammonia, a brown precipitate is formed. 
This seems to have the composition NH 3 ,NI 3 , and is named 
nitrogen iodide. It may be handled while wet. When dry, if 
touched with a feather, it decomposes into its constituents with 
violent explosion. 

Exercises. — 1. When moist air is used as a source of nitrogen, 
what advantage is there in using copper rather than the less 
expensive metal iron, for removing the oxygen (p. 60)? 

2. How many grams of water at 0° could be frozen (p. 85) by 
the removal of the heat required to evaporate 50 g. of liquid 
ammonia? 

3. How many grams of ammonia are contained in 1 1. of "con- 
centrated ammonia" (p. 342)? 

4. What are the ions of hydrazine hydrate? Formulate 
(p. 254) the neutralization of this base with sulphuric acid. 

5. What is the object attained by distilling under reduced 
pressure in making hydrazine (p. 345)? 

6. Classify (pp. 166, 258), (a) the interaction of a nitride with 
water (p. 343) and (6) of chlorine and ammonium chloride (p. 343), 
(c) the results of heating ammonium nitrite (p. 338) and (d) am- 
monium chloride (p. 345). 

7. Why does not ammonia burn in air (p. 343)? 

8. What substances are present in ammonium hydroxide 
solution? When the liquid is heated, what happens to each? 
Formulate the system. 



CHAPTER XXVI 
OXIDES AND OXYGEN ACIDS OF NITROGEN 

The names and formulae of the oxides and oxygen acids of 
nitrogen are as follows: 

Nitrous oxide N 2 < Hyponitrous acid H 2 N 2 2 

Nitric oxide NO 

Nitrous anhydride N 2 3 < > Nitrous acid HN0 2 

Nitrogen tetroxide N 2 4 and N0 2 

Nitric anhydride N 2 5 < — > Nitric acid HN0 3 . 

All the oxides are endo thermal compounds (p. 174), yet, with the 
exception of the third and the last, they are all relatively stable. 
The acids, when deprived of the elements of water, yield the oxides 
opposite which they stand (p. 281, footnote). Conversely, ex- 
cepting in the case of nitrous oxide, the anhydrides with water give 
the acids. All of these substances are made directly or indirectly 
from nitric acid — nitric anhydride by removal of water, the 
others by reduction. We turn, therefore, first, to nitric acid and 
its properties. This acid is made from Chile saltpeter (next sec- 
tion) and also by fixation of atmospheric nitrogen (see p. 352). 

Nitric Acid HN0 3 

Sources. — Sodium nitrate, or Chile saltpeter (caliche) is 
found in a desert region near the boundary of Chile and Peru. 
The deposit is about 5 feet thick, 2 miles wide, 220 miles long, and 
contains 20 to 60 per cent of the salt. Purification is effected by 
recrystallization. Potassium nitrate, or Bengal saltpeter, is found 
in the soil in the neighborhood of cities in India, Persia, and other 
oriental countries. It arises from the oxidation of animal refuse 
to nitric acid, through the mediation of nitrifying bacteria. The 
potash and lime in the soil, along with the nitric acid, give nitrates 
of potassium and calcium. The aqueous extract of this soil 
is treated with wood ashes, which contain potash K 2 C03. It is 

347 



348 COLLEGE CHEMISTKY 

poured off from the calcium carbonate thus precipitated, and is 
finally evaporated. In guano (excreta of sea birds), used as a 
fertilizer, the nitrogen compounds have often been .converted 
largely into nitrates in the same way. 

Manufacture. — When any nitrate is treated with any acid, 
nitric acid is formed by a reversible double decomposition. As 
sodium nitrate is the cheapest salt of nitric acid, it is always em- 
ployed. For the same reason and, above all, because of its relative 
involatility, sulphuric acid is used to displace it: 

NaN0 3 + H 2 S0 4 *=+ NaHS0 4 + HNOst- 

The nitric acid is rather volatile (b.-p. 86°), while sulphuric acid 
(b.-p. 330°) is much less so, and the two salts are not volatile at all. 
The materials are heated in cast-iron stills, and the nitric acid 
vapor is condensed in glass pipes surrounded by water. Thus the 
interaction proceeds to completion very easily (c/. p. 142; see also 
p. 185). 

Physical Properties. — Nitric acid is a colorless, mobile liquid 
(sp. gr. 1.52) boiling at 86°, and freezing to a solid (m.-p. —47°). 
It fumes strongly when its vapor issues into moist air (c/. p. 144). 
An aqueous solution containing 68 per cent of the acid boils at 
120.5°, while the pure acid, pure water, and all other mixtures, 
boil at lower temperatures. This 68 per cent nitric acid of constant 
boiling-point (p. 145) forms the "concentrated nitric acid" of 
commerce (sp. gr. 1.41). 

r 

Chemical Properties. — 1. Like chloric acid (p. 314), and 
other oxygen acids of the halogens, nitric acid is most stable when 
mixed with water. The pure (100 per cent) acid decomposes while 
being distilled: 

4HN0 3 -> 4N0 2 + 2H 2 + 2 , 

yet not with explosive violence like chloric acid. The distillate 
is colored brown by dissolved nitrogen tetroxide N0 2 (" fuming " 
nitric acid). Repeated distillation finally leaves 68 per cent of the 
acid, mixed with 32 per cent of water formed by the above decom- 
position. The acid of constant boiling-point is, therefore, reached, 



OXIDES AND OXYGEN ACIDS OF NITROGEN 349 

as usual, from more concentrated as well as from less concentrated 
specimens. 

2. Nitric acid, when dissolved in water, is highly ionized, and 
is therefore active as an acid. By interaction with hydroxides and 
oxides it forms nitrates. 

3. When pure nitric acid (b.-p. 86°) is poured upon phosphoric 
anhydride, the latter combines with the elements of water, and dis- 
tillation gives nitric anhydride : 2HN0 3 + P 2 5 — > N 2 5 1 + 2HP0 3 . 
The anhydride is a white solid melting at 30° and boiling at 
45°. It unites vigorously with water to form nitric acid. It 
decomposes spontaneously into nitrogen tetroxide and oxygen: 
2N 2 5 ->4N0 2 + 2 . 

4. Like the unstable oxygen acids of the halogens, nitric acid is 
an oxidizing agent even when diluted with water. The multiplicity 
of the products into which it may be decomposed by reduction, 
however, renders separate treatment of this property necessary 
(see p. 354). 

5. Nitric acid interacts energetically with many compounds of 
carbon to give nitro-derivatives. Thus, when heated with phenol 
C 6 H 5 (OH) (carbolic acid) it gives picric acid (trinitrophenol) 
C6H 2 (N0 2 )3(OH), which crystallizes in yellow needles in the mix- 
ture. This is a yellow dye, used also as an explosive. 

C6H 5 (OH) + 3HON0 2 -> C 6 H 2 (OH)(N0 2 ) 3 + 3H 2 0. 

When heated with toluene C6H 5 CH 3 , it gives trinitrotoluene: 

CH 3 C 6 H 5 + 3HON0 2 -> CH 3 C 6 H 2 (N0 2 ) 3 + 3H 2 0. 

This substance (T.N.T.) is used for filling "high explosive" shells, 
because it can be melted (m.-p. 81.5°) and poured in, making the 
filling easy, safe, rapid, and complete. It is not easily exploded by 
shocks during transportation, but it explodes instantaneously and 
completely with a detonator. The following equation shows, 
roughly, the decomposition, and the large amount of carbon set 
free explains the black smoke produced: 

2CH 3 C 6 H 2 (N0 2 ) 3 -* 5H 2 + 3N 2 + 6C0 2 + 8C. 

6. Organic compounds of another class, the alcohols (q.v.), inter- 
act with molecular nitric acid in a different way. The latter is 
mixed with sulphuric acid, which assists in the removal of the ele- 



350 COLLEGE CHEMISTRY 

ments of water (p. 286). Thus, when glycerine is added slowly to 
the cooled mixture, glyceryl nitrate (so-called nitroglycerine) is 
produced: 

C 3 H 5 (OH) 3 + 3HN0 3 -> C 3 H 5 (N03)3 + 3H 2 0. 

Guncotton is made by this action, cotton (cellulose) being em- 
ployed : 

(C 6 H 10 O 5 )2 + 6HNO3 -> C 12 H 14 4 (N0 3 )6 + 6H 2 0. 

7. Nitric acid produces substances of bright-yellow color, known 
as xanthoproteic acids, when it comes in contact with proteins, e.g., 
in the skin, or in wool. Hence nitric acid stains woolen clothing 
yellow. This reaction is used as a test for proteins. 

Nitrates. — The nitrates are all more or less easily soluble in 
water. When heated they decompose in one or other of three ways 
(see pp. 351, 356, 357). The individual nitrates, such as sodium 
nitrate and potassium nitrate, are described elsewhere. 

Nitric Oxide and Nitrogen Tetroxide 

Preparation of Nitric Oxide NO, — Pure nitric oxide is ob- 
tained by adding nitric acid to a boiling solution of ferrous sulphate 
in dilute sulphuric acid or of ferrous chloride in hydrochloric acid: 

2FeS0 4 + H 2 S0 4 -> Fe 2 (S0 4 ) 3 (+ 2H) X 3. (1) 
(3H) + HN0 3 -> NO + 2H 2 X 3. (2) 

6FeS0 4 + 3H 2 S0 4 + 2HN0 3 -* 3Fe 2 (S0 4 ) 3 + 2NO + 4H 2 0. 

The first partial equation does not take place at all unless an oxi- 
dizing agent like nitric acid is present (p. 225) . The multiplication 
of the two partial equations by 3 and 2, respectively, is required in 
order that the hydrogen, which is not a product, may cancel out. 
This action is used as a means of determining the quantity of 
nitric acid in a solution, or of nitrates in a mixture, by measure- 
ment of the volume of nitric oxide evolved. 

As we shall see (p. 354), nitric oxide may also be obtained when 
sufficiently dilute nitric acid (sp. gr. 1.2) acts upon copper. This 
interaction furnishes the most convenient method of generating 
the gas in the laboratory (see also p. 352). 



OXIDES AND OXYGEN ACIDS OF NITROGEN 351 

Properties of Nitric Oxide, — Nitric oxide is a colorless gas. 
In solid form it melts at —167°, and the liquid boils at —153.6°. 
Its solubility in water is slight. The density of the gas shows the 
formula to be NO; and there is no tendency to form a polymer, 
such as N 2 2 , even at low temperatures. 

This gas is the most stable of the oxides of nitrogen. Vig- 
orously burning phosphorus continues to burn in the gas, nitrogen 
being set free. Burning sulphur and an ignited taper, however, 
are extinguished. 

Nitric oxide has two characteristic chemical properties. It unites 
directly with oxygen in the cold to form the reddish-brown nitrogen 
tetroxide : 

2NO + 2 <f± 2N0 2 . 

The same result follows when it is led into warm concentrated 
nitric acid: NO + 2HN0 3 *± 3N0 2 + H 2 0. 

It also unites with a number of salts, the compound in the case of 
ferrous sulphate, FeNO.SO^ being capable of existence in solution 
and possessing a brown color. Since ferrous sulphate will first 
reduce nitric acid to nitric oxide (p. 350), and the excess of the salt 
will then give a brown color with the product, a delicate test for 
nitric acid is founded upon these actions. 

Preparation of Nitrogen Tetroxide N0 2 . — This substance 
is liberated by heating nitrates, other than those of potassium, 
sodium, or ammonium, such as lead and copper nitrates: 

2Cu(N0 3 ) 2 -> 2CuO + 4N0 2 + 2 . 

The oxide of the metal remains, unless this oxide is itself decom- 
posed by heating (p. 60). When the mixed gases are led through 
a U-tube immersed in ice, the tetroxide condenses as a yellow 
liquid (b.-p. 22°, m.-p. —10.5°), and the oxygen passes on. 

The compound may also be made by direct union of nitric oxide 
and oxygen, or by oxidation of nitric oxide by concentrated nitric 
acid (p. 351). It is likewise almost the sole product of the inter- 
action of concentrated nitric acid with tin or copper (see p. 355). If 
any nitric oxide were produced by the primary action, it would 
be oxidized to nitrogen tetroxide in passing up through the acid 
(p. 351). 



352 COLLEGE CHEMISTRY 

Properties of Nitrogen Tetroxide. — The most striking 
peculiarity of this gas is that, when hot, it is deep brown in color, 
and when cold, pale yellow. The density of the brown gas, at 140°, 
corresponds to the formula N0 2 , that of the yellow gas at 22°, to 
N 2 4 . When the temperature is carried above 154°, by passing the 
brown gas through a red-hot tube, the brown color disappears, and 
nitric oxide and oxygen are formed. On cooling, the same steps 
through brown gas to pale-yellow gas are retraced: 

2NO + 2 <± 2N0 2 <± N 2 4 

Colorless Brown Colorless 

Since nitrogen tetroxide yields free oxygen more readily than 
does nitric oxide, phosphorus burns readily in it; a taper, however, 
is extinguished. On account of its oxidizing power, it is sometimes 
used in bleaching flour. 

This oxide is intermediate in composition between nitrous and 
nitric anhydrides, and, when dissolved in cold water, gives both 
nitric and nitrous acids: N 2 4 + H 2 — > HN0 3 + HN0 2 . If a 
base is present, a mixture of the nitrate and nitrite of the metal is 
produced. When the water is not cooled, the nitrous acid (q.v.) f 
being unstable, gives nitric oxide and nitric acid: 3N0 2 + H 2 
<=> 2HN0 3 + NO. 

Nitric Acid prom Atmospheric Nitrogen 

The Reactions Involved, — Nitrogen and oxygen have no 
tendency to unite at room temperature to form nitric oxide. The 
union is endothermal, and is therefore favored by a high temper- 
ature (Van't Hoff's law, p. 188) : 

N 2 + 2 + 43,200 cal. <=± 2NO. 

Even at 2000°, however, using atmospheric air, only 1 per cent of 
nitric oxide is formed, and at 3000°, 5 per cent. The electric dis- 
charge actually used gives about 1 per cent. 

The mixture is next cooled, to permit the union of 2NO + 2 
+± 2N0 2 , because (p. 352) nitrogen tetroxide is decomposed at 
about 154°, and therefore cannot be formed at 2000°. 

Next, the air containing N0 2 is passed through absorbing towers, 
down which water trickles, and nitric acid is formed: 

3N0 2 + H 2 -> 2HN0 3 + NO. 



OXIDES AND OXYGEN ACIDS OF NITROGEN 



353 



The NO liberated combines with more atmospheric oxygen to 
form N0 2 , which interacts again with the water, and practically 
no nitric oxide is lost. 

Finally, the nitric acid is poured upon limestone (CaC0 3 ), and 
the calcium nitrate formed is sold for use as a fertilizer, under the 
name air saltpeter. 




k Gas * Am Oojlct 
Magnet 



Fig. 96. 



The Plant used in the Fixa- 
tion. — At Notodden and elsewhere 
watch w, x in Norway, the 

Birkeland-Eyde 
process (Fig. 96) 
is used. Hydro- 
electric power is 
employed, and an 
arc discharge be- 
tween two rods of 
carbon is spread, 
by the influence of 
large and powerful 

electromagnets, into a circular brush discharge 
several feet in diameter. The figure is a cross 
section of the space filled by the discharge, the 
small circle in the center being a section of one 
carbon rod. Air is blown through the flame in 
such a way that none can avoid passing through 
at least a part of the heated area. The yield is 
about 70 g. of nitric acid per kilowatt-hour, and 
the net earnings are $350,000 (1911). 

The Badische process, used in the same 
factories in Norway, employs a discharge 
through a tube over 20 feet long (Fig. 97). 
The stream of air rotates as it traverses the 
tube, so that every part is exposed to the dis- 
charge. The Pauling process, used at Gelsen- 
kirchen in Germany and Nitrolee, South Caro- 
lina, uses preheated air and a different arrangement of the discharge. 
For other reactions involving the fixation of atmospheric nitro- 
gen, see calcium cyanamide (q.v.) and root nodules (p. 339). 



\EnTfiAMCe 



Fig. 97. 



354 COLLEGE CHEMISTRY 

Oxidizing Actions op Nitric Acid 

When nitric acid gives up oxygen to any body, it is itself reduced. 
Hence, according to convenience, we shall refer to oxidations by, 
or reductions of nitric acid. 

Oxidation of Hydrogen. — The metals preceding hydrogen 
in the electromotive series (p. 260) displace hydrogen from nitric 
acid, as they do from other acids. With metals more active than 
zinc, such as magnesium, a great part of the hydrogen escapes in 
the free condition. But, in the case of zinc and the metals below it, 
most or all of the hydrogen is oxidized to water by the nitric acid, 
and part of the acid is reduced (see Active hydrogen, p. 360). 
Thus, with zinc and very dilute nitric acid, almost the only product, 
aside from zinc nitrate, is ammonia: 

4Zn + 8HN0 3 -> 4Zn(N0 3 ) 2 (+ 8H). (1) 

(8H) + HNO3 -> NH 3 + 3H 2 0. (2) 

NH 3 + HNO3 -> NEUNO-s. (3) 

4Zn + 10HNO 3 -^4Zn(NO 3 ) 2 + NH4NO3 + 3H 2 0. 

With the excess of nitric acid (3), ammonium nitrate is formed. 

Heavy Metals, — The less active metals, such as copper and 
silver, do not displace hydrogen from dilute acids (p. 60), but 
reduce nitric acid, nevertheless, and are converted into nitrates. 
Platinum and gold (c/. p. 287) alone are not attacked. Thus, 
copper, with somewhat diluted nitric acid (sp. gr. 1.2), gives cupric 
nitrate and nitric oxide NO. 

In making the equation for this action we may use the anhydride 
plan (p. 325), which is applicable whenever an oxygen acid gives 
an oxide by reduction. We resolve the formula of nitric acid into 
those of water and the anhydride H 2 0,N 2 5 (= 2HN0 3 ). This 
shows that the two molecules of the acid will give 2NO, and 30 
will remain: 

2HN0 3 (or H 2 0,N 2 5 ) -> H 2 + 2N0 (+ 30). (1) 

(30) + 6HNO3 + 3Cu -» 3H 2 + 3Cu(N0 3 ) 2 . (2) 

8HNO3 + 3Cu -*4H 2 + 2NO + 3Cu(N0 3 ) 2 . 

The nitric oxide is liberated as a colorless gas, but forms the brown 
tetroxide at once on meeting the oxygen of the air (p. 351). 



OXIDES AND OXYGEN ACIDS OF NITROGEN 355 

When concentrated nitric acid is used with copper, almost pure 
nitrogen tetroxide is obtained: 

2HN0 3 (or HAN2O5) -> H 2 + 2N0 2 (+ 0). (1) 

(0) + 2HN0 5 + Cu -> H 2 + Cu(N0 3 ) 2 . (2) 

4HN0 3 + Cu -> 2H 2 + 2N0 2 + Cu(N0 3 ) 2 . 

The reader should note the constant production of nitric oxide with 
diluted nitric acid, and the invariable formation of nitrogen tetroxide 
with concentrated acid. This is explained by the fact that nitrogen 
tetroxide cannot pass unchanged through a liquid containing much 
water, for it gives nitric acid and nitric oxide with the latter (p. 352) . 
Conversely, where the nitric acid is concentrated, nitric oxide, even 
if formed by the interaction with the metal, must be oxidized to 
nitrogen tetroxide as it passes up through the liquid (p. 351). 
Note, also, that the nitrate of the metal is formed, if the nitrate is 
not hydrolyzed by water, not the oxide. 

Oxidation of Non-Metals. — With non-metals the actions are 
different, in so far that these elements form no nitrates. Thus 
sulphur boiled in nitric acid gives sulphuric acid, along with nitric 
oxide, equation (3), or with nitrogen tetroxide, equation (6), or 
with both, according to the concentration of the acid (see above) : 

2HN0 8 (or H 2 0,N 2 5 ) -> 2N0 + H 2 (+ 30). (1) 

(30) + H 2 + S -> H 2 S0 4 . (2) 

2HN0 3 + S -> 2N0 + H 2 S0 4 . (3) 

2HN0 3 (or H 2 0,N 2 5 ) -» 2N0 2 + H 2 + X 3. (4) 

(30) + H 2 -h S -> H 2 S0 4 . (5) 

6HNO3 + S -> 6N0 2 + 2H 2 + H 2 S0 4 . (6) 

The reader will note that a separate equation, (3) and (6), must 
be made for the formation of each reduction product. If NO and 
N0 2 are both formed, they cannot arise from the same molecule 
of nitric acid. They result from two actions which are independent, 
although proceeding concurrently in the same vessel (cf. p. 317). 
Thus the equation: 2HN0 3 + C -> H 2 + C0 2 + NO + N0 2 , is 
a misrepresentation. It implies that equimolar quantities of the 
two oxides of nitrogen are formed. But this could occur only by 
chance, and the balance would be destroyed the next moment by 



356 COLLEGE CHEMISTRY 

the lowering in the concentration of the acid, giving the advantage 
to the nitric oxide. 

Oxidation of Compounds: Aqua Regia. — Compounds like 
hydrogen sulphide and sulphurous acid, which are easily oxidized, 
interact with nitric acid. With diluted nitric acid, the products are 
free sulphur and sulphuric acid respectively. 

The mixture of nitric acid and hydrochloric acid is known as 
aqua regia. Chlorine is set free by the oxidation of the hydro- 
chloric acid, 

HN0 3 + 3HC1 -> 2H 2 + Cl 2 + NOC1, 

and nitrosyl chloride NOC1 is also formed. The liquid thus con- 
tains several oxidizing agents, nitric acid, hypochlorous acid (from 
Cl 2 + H 2 0), and some nitrous acid. It is frequently used in 
analysis, for example to oxidize sulphur (say, in cast iron or in 
minerals), the sulphuric acid formed being estimated by precipi- 
tation and weighing of barium sulphate (p. 287). 

Aqua regia (Lat., royal water) received its name because it con- 
verted the "noble" metals, gold and platinum, into soluble com- 
pounds. This it does because the free chlorine, in presence of 
hydrochloric acid, combines to form the exceedingly stable com- 
plex ions (see pp. 505, 508) AuCLr (see chlorauric acid, and PtCl6 = , 
the negative ion of chloroplatinic acid: 

2HC1 + 2C1 2 + Pt -> H 2 PtCl 6 , or Pt + 2C1 2 + 2C1~ -> PtCl 6 =. 

Nitrous Acid, Hyponitrous Acid, and their Anhydrides 

Nitrites and Nitrous Acid. — When the nitrates of potassium 
and sodium are heated, they lose one unit of oxygen, and the 
nitrites remain: 

2NaN0 3 -> 2NaN0 2 + 2 . 

Commonly lead is stirred with the melted nitrate and assists in the 
removal of the oxygen. The litharge PbO which is formed re- 
mains as a residue when the sodium nitrite is dissolved for re- 
crystallization. 

When an acid is added to a dilute solution of a nitrite, a pale-blue 
solution containing nitrous acid HN0 2 is obtained. The acid is 



OXIDES AND OXYGEN ACIDS OF NITROGEN 357 

very unstable, however, and, when the solution is warmed, it de- 
composes : 

3HN0 2 -» HN0 3 + 2NO + H 2 0. 

When a concentrated solution of sodium nitrite (or the solid salt 
itself) is acidified, the nitrous acid decomposes at once, and a brown 
gas containing the anhydride escapes: 

2H+ + 2N0 2 ~ U 2HN0 2 fcj H 2 + N 2 3 T . 

This behavior distinguishes a nitrite from a nitrate. 
Nitrous acid is an active oxidizing agent : 

2HI + 2HN0 2 (or H 2 0,N 2 3 ) -> 2H 2 + 2NO + I 2 . 

Indigo is also converted by it into isatin (cf. p. 311). 
Nitrous acid is much used in the making of organic dyes. 

Nitrous Anhydride N 2 3 , — A study of the density of the gas 
arising from the decomposition of nitrous acid shows that, in the 
gaseous state, the anhydride is almost entirely dissociated: 

N 2 3 *± NO + N0 2 . 

When the mixture is led through a U-tube immersed in a freezing 
mixture at —21°, a deep-blue liquid is obtained which is the 
anhydride itself. This dissociates rapidly when allowed to boil. 

The same equi molar mixture of the two gases is obtained by the 
action of water on nitrosylsulphuric acid (p. 281). 

Hyponitrous Acid and Nitrous Oxide N 2 0. — Hyponitrous 
acid H 2 N 2 2 is a white solid. Its solution in water is an exceed- 
ingly feeble acid. The warm aqueous solution decomposes slowly, 
giving nitrous oxide : 

H 2 N 2 2 -> H 2 + N 2 0, 

and this change is not capable of reversal. 

Nitrous oxide is prepared by gently heating ammonium nitrate 
(an explosive), or a solution of a salt of ammonium and a nitrate; 

NH4+ + N0 3 " <± NH4NO3 -> 2H 2 + N 2 0. 

The steam condenses, and the nitrous oxide may be collected over 
warm water, or be dried and compressed into steel cylinders. 



358 COLLEGE CHEMISTRY 

Its solubility in cold water is considerable: at 0°, 130 volumes in 
100; at 25°, 60 in 100. The liquefied gas boils at -89.8° and its 
vapor tension at 20° is 49.4 atmospheres. 

A glowing splinter of wood bursts into flame in nitrous oxide, 
and phosphorus and sulphur burn in it with much the same vigor 
as in oxygen. In all cases oxides are formed, and nitrogen is set 
free. It does not combine with nitric oxide, however, as does 
oxygen (p. 351). 

Metals do not rust in nitrous oxide, and the haemoglobin of the 
blood is unable to use it as a source of oxygen. It is employed as an 
anaesthetic for minor operations. The hysterical symptoms which 
accompany its use caused it to receive the name of " laughing gas." 

Graphic Formulae of Nitric Acid and its Derivatives: Ex- 
plosives. — The following equation for the formation of am- 
monium nitrate by neutralization of ammonium hydroxide with 
nitric acid shows the graphic (p. 292) or structural formulae of 
these substances: 

H \ O H \ O 

S x N-OH+H-0-Nf -> S^N-O-Nf +H 2 0. 
H / O H / O 

The structural formula of the nitrate is intended to explain the fact 
that the salt is able to exist at all, by representing the oxygen and 
hydrogen as being separated from one another and attached to 
different nitrogen units. When the equilibrium of the system is 
disturbed by heating, the oxygen and hydrogen unite to form 
water, an arrangement which is much more stable, and nitrous 
oxide (p. 357) escapes with the steam. 

The behavior of nitroglycerine and guncotton (p. 350), as well 
as of ammonium nitrite (p. 338), is explained in the same way. 
These substances are made by actions which, like the above 
neutralization, take place in the cold, and the groups, containing the 
oxygen on the one hand and carbon and hydrogen on the other, 
become quietly united without more serious interaction. When, 
however, the nitroglycerine, for example, is heated, or receives a 
mechanical shock, the carbon and hydrogen unite with the oxygen 
and the nitrogen escapes: 

4C 3 H 5 (N0 3 )3 -> 12C0 2 + 10H 2 O + 6N 2 + 2 . 



OXIDES AND OXYGEN ACIDS OF NITROGEN 359 

Smokeless Powder and Dynamite. — Dried guncotton (p. 
350) simply burns briskly (deflagrates) when set on fire. Whether 
wet or dry, it explodes, but only from a suitable shock, such as 
that produced by fulminate of mercury Hg(OCN) 2 , used in per- 
cussion caps. In pure form it is used only in torpedoes or sub- 
marine mines. Like nitroglycerine (p. 350), it explodes too 
rapidly, and would burst the gun, or pulverize the ore or coal if 
used for blasting. Neither of these substances " explodes down- 
wards only." The explosion strikes the air with equal violence, 
but the effect on the air escapes notice because it is not permanent, 
while the shattering of a rock or plate of steel remains. 

Cordite, one variety of smokeless powder, is made by dissolving 
guncotton (65 parts), nitroglycerine (30 parts) and vaseline (5 
parts) in acetone. The resulting paste is rolled out and cut into 
small pieces. When the acetone evaporates, the horny cordite 
remains. These explosives are smokeless because, unlike gun- 
powder and T.N.T., they yield no solids when they decompose 
(see equations). 

Various forms of dynamite are made like cordite, excepting that 
sodium or ammonium nitrate and sawdust or flour are added, so 
that the rate of explosion may be regulated and the coal or ore may 
be split up, but not shattered or pulverized. 

Plastics. — A guncotton, less completely " nitrated" by nitric 
acid, when worked between rollers Yvith camphor and a little 
alcohol, gives a viscous solution (Parkes, before 1855). When the 
alcohol evaporates, transparent, colorless celluloid (first made by 
Hyatt) remains. The moist dough is rolled into sheets to make 
photographic films. By adding dyes and "fillers," and molding 
the dough, black combs, brush handles, white knife handles, etc., 
can be manufactured. 

Collodion is a solution of the same guncotton in a mixture of 
alcohol and ether. When collodion is forced through minute holes 
in a steel dye, the threads dry as they come out and can be wound 
on spools. Treatment with an alkali " denitrates" the threads, 
restoring the composition to that of the original cotton. The prod- 
uct, one of the forms of artificial silk, is at least as brilliant as the 
real article (a protein, not related chemically to cotton), and sus- 
ceptible of being dyed to any desired tint. 



360 COLLEGE CHEMISTRY 

Balancing Equations. — The reader should practice the 
balancing of the equations for oxidations occurring in this chapter, 
using all the methods. We have used the anhydride plan (p. 355) 
#nd that of partial equations (p. 350). To illustrate the other 
plans, take the action of concentrated nitric acid on copper (p. 
355). 

Positive and negative valence method (p. 322). Write the skeleton 
equation : 

Skeleton: HN0 3 + Cu -> H 2 + N0 2 + Cu(N0 3 ) 2 . 

We perceive that on the left the valence of 3 is —6 and of H is 
+ 1 : that of N is therefore +5. That of Cu is zero. On the right, 
the valence of N is +4 and of Cu +2. Evidently, 2N changing 
from 2 X +5 to 2 X +4 will furnish +2 for the copper. We 
note also that 2N0 3 is required, without change, for Cu(N0 3 ) 2 . 
Hence, altogether 4HN0 3 is needed on the left, and gives 2N0 2 : 

Balanced: 4HN0 3 + Cu -> 2H 2 + 2N0 2 + Cu(N0 3 ) 2 . 

Positive electrical charge plan (p. 325). In the skeleton equation 
(above) we first separate the oxidizing ions and their products from 
the oxidized substance and the change it undergoes: 

NOr + 2H+ -> N0 2 ° + H 2 + X 3. 

Cu° + 20 -^Cu++. 

Cu° + 2N0 3 " + 4H+ -» 2N0 2 ° + 2H 2 + Cu++. 

The first partial equation produces ©, while the second requires 
2 ©, and hence the former is multiplied by 2 before the addition 
takes place. Since N0 3 ~ is the only acid radical present, it is 
understood that cupric nitrate is the salt formed. 

Active ("Nascent") Hydrogen. — When hydrogen gas is led 
through cold nitric acid, little or no action occurs. But (p. 354) 
when zinc, which interacts with acids to give hydrogen, is placed 
in nitric acid the latter is reduced. To explain the apparent 
greater activity of the hydrogen in the second instance, we note 
the fact that it is liberated on the surface of the zinc. The contact 
(catalytic) effect of the zinc increases its activity. Many metals 
have, in a greater or less degree, this power of increasing the 
activity of hydrogen. Thus, hydrogen absorbed in platinum or 



OXIDES AND OXYGEN ACIDS OF NITROGEN 361 

palladium (p. 57) or liberated by electrolysis on poles made of 
these metals, reduces nitric acid readily. Other elements, such as 
the oxygen used in making sulphur trioxide (p. 279), are also ren- 
dered more active by contact agents. 

This more active state of hydrogen is described as the nascent 
state, because it happens to be a common condition of hydrogen 
when associated with substances which produce it. The active 
state has, however, no necessary connection with such an immedi- 
ately preceding act of liberation, as the platinum and sulphur 
trioxide illustrations, and the following experiment show: Three 
test-tubes are filled with dilute, acidified potassium permanganate 
solution. Zinc dust, added to one, generates hydrogen and causes 
decolorization. A little platinum black is added to the second, and 
hydrogen gas is led through this and the third. The contact 
action of the platinum enables the hydrogen quickly to reduce 
the permanganate, while the third portion remains unaltered. 

Besides, if the action were due to freshly liberated, perhaps 
atomic hydrogen, this substance should have constant properties. 
But it has not. Thus, nitric acid with zinc gives much ammonia; 
with magnesium, none; with tin, ammonia and hydroxylamine 
HONH 2 as weU. 

Exercises. — 1. Make the equation for the interaction of 
ferrous chloride, hydrochloric acid, and nitric acid (p. 350), and 
for all the actions concerned when the test for a nitrate (p. 351) is 
applied to sodium nitrate. What volume (at 0° and 760 mm.) of 
NO is obtained from one formula-weight of nitric acid (p. 350)? 

2. In the action of zinc on dilute nitric acid (p. 354), why is not 
the ammonia given off as a gas? How should you show that it was 
formed at all? 

3. Make correct equations for the formation of nitric oxide and 
nitrogen tetroxide by the action of carbon on nitric acid (p. 355). 

4. Make equations for the interaction of iron with diluted and 
with concentrated nitric acid, respectively (p. 355) . The iron gives 
ferric nitrate Fe(N0 3 )3- 

5. Give the three ways in which nitrates decompose when 
heated, with one equation illustrating each. 

6. Make all the equations for oxidations on pp. 350 and 354, 
using the methods illustrated on p. 360. 



CHAPTER XXVII 

PHOSPHORUS 

The Chemical Relations of the Element. — There are 
many things in the chemistry of phosphorus and its compounds 
which remind us of nitrogen. Yet these are largely referable to 
the fact that the elements are both non-metals and both have the 
same valences, viz., three and five. The behavior of the com- 
pounds is often very different. For the present it is sufficient to 
say that both give compounds with hydrogen, NH 3 and PH 3 , and 
both yield oxides of the forms X 2 3 , X2O4, and X 2 5 . The first 
and last of these oxides are acid-forming, and phosphorus, there- 
fore, gives acids corresponding to nitrous and nitric acids. The 
element is thus a non-metal. 

Occurrence. — This element is found in nature in the form of 
phosphates. Calcium phosphate Ca 3 (PC>4)2 forms 26 per cent of 
the bones and teeth, and it occurs in all fertile soils. It consti- 
tutes a large part of the " phosphate rock" of Georgia, Florida, 
the Carolinas, Tennessee, and of Algeria and Tunis. A con- 
spicuous mineral related to this substance, apatite, Ca 5 F(P0 4 ) 3 
and Ca 5 Cl(P0 4 ) 3 , is found in large quantities in Canada, and is 
a component of many rocks. Complex organic compounds of 
phosphorus, such as lecithin, are essential constituents of the 
muscles, the nerves, and the brain. Amongst foods, egg-yolks and 
beans contain a large proportion. 

Preparation. — Brand, merchant and alchemist, of Hamburg, 
discovered phosphorus (1669) by distilling the residue from evapo- 
rated urine, in the course of his search for the philosopher's stone. 
The mode of preparing it from bone-ash, which contains 83 per 
cent of calcium phosphate, was first published by Scheele (1771). 
Now the less expensive calcium phosphate of fossil origin is 
employed. 

362 



PHOSPHORUS 



363 



The calcium phosphate is mixed with the proper proportions of 
carbon and silicon dioxide (sand), and the mixture is introduced 
continuously into an electric furnace (Fig. 98) . The discharge of 
an alternating current between carbon poles 
produces the very high temperature which 
the action requires. The calcium silicate 
which is formed fuses to a slag, and can be 
withdrawn at intervals. The gaseous prod- 
ucts pass off through a pipe and the phos- 
phorus is caught under water: 

Ca 3 (P0 4 ) 2 + 3Si0 2 + 5C -> 3CaSi0 3 
+ 5CO + 2P. 




Fig. 98. 



We may regard the phosphate as being com- 
posed of two oxides, 3CaO,P 2 5 . It thus 
appears that the calcium oxide has united 
with the silica, which is an acid anhydride (c/. p. 280) : CaO + Si0 2 
— * CaSi0 3 , while the phosphoric anhydride has been reduced. 

The phosphorus, after purification, is cast into sticks in tubes of 
tin or glass, standing in cold water. 

The Electric Furnace. — By an electric furnace is understood 
an electrothermal arrangement in which the heat produced by some 
resistance offered to the current, such as that of an air-gap 
between the carbons, is used to produce chemical changes re- 
quiring a high temperature. Electrolysis plays no part in the 
phenomena, and an alternating current, which can produce no 
electrolytic decomposition, is generally employed. The restricted 
area within which the heat is developed makes possible the attain- 
ment of a high temperature (see Calcium carbide). 



Physical Properties. — There are at least two allotropic forms 
(p. 222) of phosphorus, known as white phosphorus and red 
phosphorus. White phosphorus, prepared as described above, is at 
first transparent and colorless, but after exposure to light acquires 
a superficial coating of the red variety. It melts at 44° and boils 
at 287°. Its sp. gr. is 1.83. Its molecular weight at 313° is 128 
and the formula, therefore, P4 (cf. p. 111). Yellow phosphorus 
is soluble in carbon bisulphide, less soluble in ether, and insoluble 



364 COLLEGE CHEMISTRY 

in water. It is exceedingly poisonous, less than 0.15 g. being a 
fatal dose, and is an ingredient in roach-paste and rat poison. 
Continued exposure to its vapor causes necrosis, a disease from 
which match-makers are liable to suffer. The jawbones and teeth 
are particularly liable to attack. 

Red phosphorus is a red powder consisting of small tabular 
crystals. It is obtained by heating yellow phosphorus to about 
250° in a vessel from which air is excluded. Much heat is evolved 
in the transformation. Red phosphorus does not melt, but 
passes directly into vapor, identical with that of yellow phos- 
phorus. It is insoluble in carbon bisulphide and other solvents. 
It is not poisonous, and, unlike yellow phosphorus, does not re- 
quire to be kept under water to avoid spontaneous combustion. 
Red phosphorus appears to be a solid solution (p. 122) of the 
white variety in a less active kind. Hence, its properties are 
variable, e.g., sp. gr. from 2.05 to 2.34. Bridgeman, by heating 
white phosphorus at 200° under a pressure of 1200 kg./cm 2 ., has 
obtained black phosphorus (sp. gr. 2.69) which may be the pure 
form of the red variety. 

Chemical Properties. — White phosphorus unites directly 
with the halogens with great vigor. It unites slowly with oxygen 
in the cold, and with sulphur and many metals when the materials 
are heated together. The slow union of cold phosphorus with 
atmospheric oxygen is accompanied by the evolution of light. 
Hence the word phosphorescence. The name of the element (Gk. 
<f>u>s, light; <£e/3w, I bear) records this property. Apparently the 
chemical energy, transformed in connection with the oxidation, 
is converted, in part at least, into radiant energy instead of com- 
pletely into heat.* The slow oxidation of phosphorus is ac- 
companied by the production of ozone, but the nature of the 
action is still unknown (cf. p. 219). 

Red phosphorus, since it is formed with evolution of heat, con- 
tains less energy than white phosphorus and is much less active. 
It does not catch fire below 240°, while ordinary phosphorus 
ignites at 35°. 

* The same production of light from chemical action in a cold body is 
seen in the luminosity of certain parts of fireflies and some species of 
fish. 



PHOSPHORUS 365 

Matches. — In making common matches, invented in 1827, 
the splints are first dipped in melted sulphur or paraffin to the 
extent of about half an inch. The head is often composed of 
lead dioxide Pb0 2 , which supplies oxygen, a small proportion of 
free phosphorus or a sulphide of phosphorus P 4 S 3 which is readily 
ignited by friction, and dextrin or glue. The use of white phos- 
phorus is forbidden by law in Sweden, France, Great Britain and 
Switzerland, and is prevented by a tax of two cents per 100 
matches in the United States. 

In the case of "safety" matches, the mixture upon the head is 
not easily ignited by itself. It is composed of potassium chlorate 
or dichromate, some sulphur or antimony trisulphide Sb 2 S3 (com- 
bustible), and a little powdered glass or chalk to increase the 
friction, all held together with glue. Upon the rubbing surface 
on the box is a thin layer of antimony trisulphide mixed with red 
phosphorus, chalk or glass, and glue. The friction converts a 
little of the red phosphorus into vapor, which catches fire readily. 

Phosphine. — Three hydrides of phosphorus are known. 
These are, phosphine PH 3 (a gas), a liquid hydride P2H4, which is 
presumably the analogue of hydrazine (N2H4), and a solid hydride 
P 4 H 2 . 

Phosphine PH 3 is formed slowly by the action of active hydrogen, 
from zinc and hydrochloric acid at 70°, upon white phosphorus. 
The gas may be made by boiling white phosphorus with potassium 
hydroxide solution in a flask provided with a delivery tube. 
Potassium hypophosphite is formed at the same time: 

3KOH + 4P + 3H 2 -> 3KH 2 P0 2 + PH 3 T . 

The gas made in this way contains a little of the vapor of the liquid 
hydride, which is spontaneously inflammable, and consequently the 
bubbles of the mixture catch fire when they reach the surface of 
water in the trough: PH 3 + 20 2 — > H3PO4. In still, moist air, 
the fog of droplets of phosphoric acid solution form smoke 
rings. 

The simplest method of preparing phosphine is by the action of 
water upon calcium phosphide: 

Ca 3 P 2 + 6H 2 -> 3Ca(OH) 2 + 2PH 3 . 



366 COLLEGE CHEMISTRY 

This action is analogous to that of water upon magnesium nitride 
(p. 339), by which ammonia is produced. In consequence of the 
fact that calcium phosphide is a substance of irregular compo- 
sition, a mixture of all three hydrides is generally obtained. By 
passing the gas through a strongly cooled delivery tube, however, 
the liquid and solid compounds are condensed and fairly pure 
phosphine passes on. 

Phosphine is a colorless gas, which is easily decomposed by heat 
into its elements. It is exceedingly poisonous and, unlike am- 
monia, it is insoluble in water, and produces no basic compound 
corresponding to ammonium hydroxide when brought in contact 
with this substance. It resembles ammonia, formally at least, in 
uniting with the hydrogen halides (see below). It differs from 
ammonia, however, inasmuch as it does not unite with the oxygen 
acids. Phosphine acts upon solutions of some salts, precipitating 
phosphides of the metals: 

3CuS0 4 + 2PH 3 -> Cu 3 P 2 1 + 3H 2 S0 4 . 

Phosphonium Compounds. — Hydrogen iodide unites with 
phosphine to form a colorless solid, crystallizing in beautiful, 
highly refracting, square prisms: PH 3 + HI — -> PELJ. Hydrogen 
chloride combines similarly with phosphine, but only when the 
gases are cooled by a freezing mixture, or are brought together 
under a total pressure of 18 atmospheres at 14°. When the 
pressure is released, rapid dissociation occurs. 

In imitation of the ammonia nomenclature, these substances 
are called phosphonium iodide and phosphonium chloride PH4CI. 
They are entirely different, however, from the corresponding am- 
monium derivatives, for the PH4+ ion is unstable. When brought 
in contact with water they decompose into their constituents, the 
hydrogen halide going into solution, and the phosphine being 
liberated as a gas. 

Halides of Phosphorus. — The existence of the following 
halides has been proved conclusively: 

.... P2I4 (solid) 

PF 3 (gas) PCI3 (liquid) PBr 3 (liquid) PI 3 (solid) 
PF 6 (gas) PCls (solid) PBr 5 (solid) 






PHOSPHORUS 367 

These substances may all be formed by direct union of the elements. 
They are incomparably more stable than are the similar com- 
pounds of nitrogen. They are all hydrolyzed by water, and give 
an oxygen acid of phosphorus and the hydrogen halide (see below). 
This action was used in the preparation of hydrogen bromide 
(p. 197) and hydrogen iodide (p. 201). 

Phosphorus trichloride PC1 3 is made by passing chlorine gas over 
melted phosphorus in a flask until the proper gain in weight has 
occurred. The substance, which is a liquid boiling at 76°, is stable 
(cf. p. 93). When excess of chlorine is employed, phosphorus pen- 
tachloride PCI5, which is a white solid body, is formed. When 
moist air is blown over any of these substances, the water is con- 
densed to a fog by the hydrogen halide. In the case of the inter- 
action of phosphorus pentachloride and water, phosphoric acid is 
formed : 

PCI5 + 4H 2 -> H 3 P0 4 + 5HC1. 

Phosphorus pentachloride, when heated, reaches a vapor 
tension of 760 mm. at 163°, and while still solid. At higher 
pressure it melts at 166°. At these temperatures, about 4 per cent 
of the molecules are dissociated into phosphorus trichloride and 
chlorine (p. 117) : PC1 5 <=± PC1 3 + Cl 2 . 

Oxides of Phosphorus. — The oxides of phosphorus are the 
so-called trioxide P4O6, the pentoxide P2O5, and a tetroxide P2O4. 

The pentoxide is a white powder formed when phosphorus is 
burned with a free supply of oxygen. It unites with water with 
great violence to form metaphosphoric acid (see below), and hence 
is known as phosphoric anhydride : P 2 5 + H 2 — » 2HPO3. In the 
laboratory this action is frequently utilized for drying gases 
(p. 330) and for removing water from combination (p. 349). The 
vapor density leads to the formula P4O10, use of which, however, 
would only complicate our equations. 

The trioxide P 4 Oe is obtained by burning phosphorus in a tube 
with a restricted supply of air. It is a white solid, melting at 
22.5° and boiling at 173°. This oxide is the anhydride of phos- 
phorous acid, but it unites exceedingly slowly with cold water to 
form this substance. It interacts vigorously with hot water, but 
phosphine, red phosphorus, hypophosphoric acid, and phosphoric 
acid are amongst the products, and very little phosphorous acid 



368 COLLEGE CHEMISTRY 

escapes decomposition. When this oxide is heated to 440° it de- 
composes, giving the tetroxide P2O4 and red phosphorus. 

Acids of Phosphorus. — There are six different acids of phos- 
phorus. Three are phosphoric acids, representing the same stage 
of oxidation of phosphorus, but different degrees of hydration of 
the anhydride. The others show three different and lower states 
of oxidation (study by positive and negative valences, p. 323) : 

Orthophosphoric acid H3PO4 ( = 3H 2 0,P 2 6 ) 
Pyrophosphoric acid H4P2O7 (= 2H 2 0,P 2 6 ) 
Metaphosphoric acid HP0 3 ( = H 2 0,P 2 5 ) 
Hype-phosphoric acid H 2 P0 3 ( = 2H 2 ; P 2 4 ) 
Phosphorous acid H 3 P0 3 (= 3H 2 0,P 2 3 ) 

Hypophosphorous acid H 3 P0 2 ( = 3H 2 0,P 2 0) 

The Phosphoric Acids. — The relation between the three 
different phosphoric acids may be seen by considering them as 
being formed from phosphorus pentoxide (the anhydride) and 
water. In the majority of cases already considered this sort of 
action takes place in but one way. Thus, nitric acid is known in 
but one form, which is produced by the union of one molecule 
each of nitrogen pentoxide and water: N 2 5 + H 2 — > 2HN0 3 . 
Similarly, the chief sulphuric acid is the one formed from one 
molecule of sulphur trioxide and one molecule of water: SO3 + 
H 2 — » H2SO4, although here we have also disulphuric acid 
H 2 S 2 7 , or H 2 0,2S0 3 . 

Now, when phosphoric anhydride acts upon water we obtain a 
solution which, on immediate evaporation, leaves a glassy solid, 
HPO3, known as metaphosphoric acid. This is H 2 0,P 2 5 . When, 
however, the solution is allowed to stand for some days, or is 
boiled with a little dilute nitric acid, whose hydrogen-ion acts 
catalytically, the residue from evaporation is H3PO4, orthophos- 
phoric acid: 

P 2 5 + 3H 2 0->2H 3 P0 4 or HP0 3 + H 2 -» H 3 P0 4 . 

This acid is 3H 2 0,P 2 0s, and no further addition of water to form 
a different acid (see p. 370) can be effected. 

Conversely, when orthophosphoric acid is kept at about 255° 
for a time, it slowly loses water, and H4P 2 07, pyrophosphoric acid, 
is obtained: 2H3 pQ 4 ^ g^ + EJp ^ 



PHOSPHORUS 



369 



This acid is 2H 2 0,P 2 5 . Further desiccation leaves metaphos- 
phoric acid, which cannot be further resolved into phosphorus 
pent oxide and water. When dissolved in water, pyrophosphoric 
acid slowly resumes the water which it has lost and gives the ortho- 
acid again. The relations of all these substances are more clearly 
seen in the graphic formulae: 



= 

-0-H 
-0-H 
-0-H 

-0-H 
-0-H 
-0-H 

= 





0-H 

0-H 



0-H 

0-H 





P - 












0- 


-H 




= 







P< 


= 
















»l 


= 


0- 


-H 


. = 












A most important fact to be noted is that the addition or removal of 
water leaves the valence of the phosphorus unchanged. The degree of 
oxidation of the phosphorus and its valence are identical in the 
three acids. 



The Relations of Anhydrides and Oxygen Acids. — Con- 
sidering the anhydride from which an acid is derived gives us the 
key to the nature and behavior of the acid. It tells much that 
the formula of the acid does not tell. For example: (1) What is 
the valence of phosphorus in H3PO3? The only way to answer 
this question is to resolve the formula (doubled, if necessary) into 
water and the anhydride, 3H 2 0,P 2 3 . The phosphorus is triva- 
lent. (2) How is this acid related to metaphosphoric acid HP0 3 ? 
Resolve the latter, as before, H 2 0,P 2 5 . The answer is that in the 
latter the phosphorus is quinquivalent. (3) How can we get 
phosphorous acid from metaphosphoric acid, or vice versa? Con- 
sidering the anhydrides, we answer, by reduction and oxidation, 
respectively. (4) Is pyrophosphoric acid H4P2O7, because it con- 
tains 70, to be made from all others by oxidation? Resolve it into 
water and anhydride, 2H 2 0,P 2 5 . We then perceive that to make 
it from phosphorous acid requires oxidation, but to make it from 
ortho- or metaphosphoric acid requires only a change in the de- 
gree of hydration: adding or subtracting water, since it adds or 
subtracts hydrogen and oxygen in equivalent amounts, is not 



370 COLLEGE CHEMISTRY 

oxidation or reduction. These conceptions have been discussed be- 
fore (pp. 316, 321). 

Orthophosphoric Acid HqPO^, — The impure, commercial acid 
is made by mixing selected, pulverized phosphate rock Ca 3 (P0 4 ) 2 
with sulphuric acid (sp. gr. 1.5) and heating with steam and 
stirring in a wooden vat. 

Ca 3 (P0 4 ) 2 + 3H 2 S0 4 £=► 2H 3 P0 4 + 3CaS04 . 

The calcium sulphate is precipitated during the heating and the 
subsequent concentration of the nitrate. 

Pure orthophosphoric acid may be made by boiling red phos- 
phorus with slightly diluted nitric acid and evaporating the water 
and excess of nitric acid. The product of recrystallization is a 
white, crystalline, deliquescent, solid hydrate, 2H 3 P0 4 ,H 2 0. 

The acid is much weaker than sulphuric acid, and is dissociated 
chiefly into the ions H+ and H 2 P0 4 ~. The dihydrophosphate-ion 
H 2 P0 4 ~~, being an acid as well as an ion, is further broken up to some 
extent into H + and HP0 4 = , as we learn from the fact that the 
solution of the sodium salt NaH 2 P0 4 is acid. The ion HP0 4 = is 
hardly dissociated at all, for a solution of the salt Na 2 HP0 4 is 
not acid in reaction. 

Salts of Orthophosphoric Acid. — As a tribasic acid, it forms 
salts of three kinds, such as NaH 2 P0 4 , Na 2 HP0 4 , and Na 3 P0 4 . 
These are known respectively as primary, secondary, and tertiary 
sodium orthophosphate. The primary sodium phosphate is 
faintly acid in reaction. The secondary one is slightly alkaline, 
because of hydrolysis arising from the tendency of the hydrogen- 
ion of the water to combine with the HP0 4 = to form H 2 P0 4 ~, 
which is much more feebly acid than is phosphoric acid H 3 P0 4 . 
The simplified equation (p. 271) shows the reason for the alka- 
linity of the solution: HP0 4 = + H+ + OH - -> H 2 P0 4 ~ + OH~ 
for hydroxyl-ion is present. The tertiary phosphate is stable only 
in solid form, and can be made by evaporating to dryness a 
mixture of the secondary phosphate and sodium hydroxide: 
Na 2 HP0 4 + NaOH <± Na 3 P0 4 + H 2 f . 

When the product is dissolved in water, this action is reversed (c/. 
p. 271). Mixed phosphates are also known, particularly sodium- 



PHOSPHORUS 371 

ammonium phosphate (microcosmic salt) NaNH4HP0 4 ,4H 2 0, 
and the insoluble magnesium-ammonium phosphate MgNKtPC^. 
Primary calcium phosphate (q.v.), known in commerce as "super- 
phosphate," is used as a fertilizer. 

The tertiary phosphates are unchanged by heating. The primary 
and secondary phosphates, however, retaining, as they do, some of 
the original hydrogen of the phosphoric acid, are capable of losing 
water, like phosphoric acid itself, when heated. The actions are 
slowly reversed when the products are dissolved in water: 

NaH 2 P0 4 <=»NaP0 3 + H 2 OT. 
2Na 2 HP0 4 <=* Na 4 P 2 7 + H 2 f . 

It will be seen that the meta- and pyrophosphates of sodium are 
formed by these actions; and this is indeed the simplest way of 
forming these substances, since the acids themselves are not per- 
manent in solution, and are too feeble to lend themselves to exact 
neutralization. Ammonium salts of phosphoric acid lose am- 
monia, as well as water, when heated (cf. p. 344, last par.). Thus, 
microcosmic salt gives primary sodium phosphate: 

NaNI^HPO* -» NH 3 1 + NaH 2 P0 4 -> NaP0 3 + H 2 1 , 

and this in turn is converted into the metaphosphate by loss of 
water. 

Pyrophosphoric Acid and Metaphosphoric Acid. — Pyro- 

phosphoric acid HJ^O?, although tetrabasic, gives only the neutral 
salts, such as Na 4 P 2 7 , and those in which one-half of the hydrogen 
has been displaced by a metal, such as Na 2 H 2 P 2 7 . 

Metaphosphoric acid HPO3 is the " glacial phosphoric acid" of 
commerce, and is usually sold in the form of transparent sticks. It 
is obtained by heating orthophosphoric acid, or by direct union of 
phosphorus pentoxide with a small amount of cold water. It 
passes into vapor at a high temperature, and its vapor density 
corresponds to the formula (HP0 3 ) 2 . 

Sodium metaphosphate NaP0 3 , in the form of a small globule 
obtained by heating microcosmic salt on a platinum wire, is used 
in analysis. When minute traces of oxides of certain metals are 
placed upon such a globule, known as a bead, and heated in the 
Bunsen flame, the mass is colored in various tints according to the 



372 COLLEGE CHEMISTRY 

oxide used (bead test). This action may be understood when we 
consider that sodium metaphosphate takes up water to form 
primary sodium orthophosphate : NaP0 3 + H 2 — > NaH 2 P0 4 . 
In the same way, but at higher temperatures, it is able to take up 
oxides of elements other than hydrogen, giving mixed ortho- 
phosphates. Thus, with oxide of cobalt a part of the metaphos- 
phate unites according to the equation: 

NaP0 3 + CoO -> NaCoP0 4 , 
and the product gives a blue color to the bead. 

Distinguishing Tests. — When a solution of nitrate of silver is 

added to a solution of orthophosphoric acid, or to any soluble 
orthophosphate, a yellow precipitate of silver orthophosphate 
Ag 3 P04 is produced. This is a test for orthophosphate-ion. With 
pyrophosphoric acid or any pyrophosphate the product is white 
Ag 4 P 2 7 . With metaphosphoric acid a white precipitate, AgP0 3 , 
is obtained also. Metaphosphoric acid coagulates a clear solu- 
tion (colloidal suspension) of albumin (say, white of egg), while 
ortho- or pyrophosphoric acid has no visible effect upon it (p. 417). 

Phosphorous Acid H3PO3. — When added to cold water, phos- 
phorus trioxide (P4O6) yields phosphorous acid very slowly. With 
hot water the action is exceedingly violent and complex (p. 367). 
This acid may be obtained easily by the action of water upon 
phosphorus trichloride, tribromide (p. 197), or tri-iodide and 
evaporation of the solution : 

PC1 3 + 3H 2 -> P(OH) 3 + 3HC1 1 . 

Some of this acid, along with phosphoric acid and hypophosphoric 
acid, is formed when moist phosphorus oxidizes in the air. 

In spite of the presence of three hydrogen atoms, this acid is 
dibasic, and two only are replaceable by metals. To express this 
fact, the first of the following formulae is preferred: 





since the symmetrical formula would indicate no difference be- 
tween the three hydrogen atoms. H united directly to P, as 



PHOSPHORUS 373 

here and in PH 3 , is not acidic. Phosphorous acid is a powerful 
reducing agent, precipitating silver, for example, in the metallic 
form from solutions of its salts. When heated, it decomposes, 
giving the most stable acid of phosphorus (cf. pp. 290, 308, 314, 
320, 357), namely, metaphosphoric acid, and phosphine: 

4H 3 P0 3 -> 3HP0 3 + 3H 2 + PH 3 . 

Sulphides of Phosphorus. — White phosphorus, when heated 
with sulphur, unites with explosive violence. By using red phos- 
phorus the action can be controlled. By employing the proper 
proportions the pentasulphide P 2 S 5 is secured. It is purified by 
distillation, and is a yellow crystalline solid (m.-p. 274°, b.-p. 
530°). Phosphorus pentasulphide is hydrolyzed by cold water: 

P 2 S 5 + 8H 2 -> 2H3PO4 + 5H 2 S. 

Other sulphides, P 4 S 3 (used in making matches), P 2 S 3 , and 
P 3 S 6 , may be prepared by using the constituents in the proportions 
represented by these formulae. 

Comparison of Phosphorus with Nitrogen and with Sul- 
phur. — Although phosphorus and nitrogen are regarded as be- 
longing to one family, the differences between them are more 
conspicuous than the resemblances. The latter are confined al- 
most wholly to matters concerned with valence. The differences 
are seen in the facts that nitrogen is a gas and exists in but one 
form, while phosphorus is a solid occurring in two varieties, and 
that the former is inactive and the latter active. The contrasts 
between phosphine and ammonia (p. 366) and between the halides 
of the two elements (pp. 346, 367) have been noted already. 
The pentoxide of nitrogen decomposes spontaneously; that of 
phosphorus is one of the most stable of compounds. Nitric acid 
is very active, both as acid and oxidizing agent; the phosphoric 
acids are quite the reverse. 

On the other hand, the resemblance of phosphorus to sulphur 
is marked. Both are solids, existing in several forms. Both 
yield stable compounds with oxygen and chlorine. The hydrogen 
compounds interact with salts to give phosphides of metals and 
sulphides of metals, respectively. Against these must be set the 
facts, that hydrogen sulphide does not unite with the hydrogen 



374 COLLEGE CHEMISTRY 

halides at all while phosphine gives the phosphonium halides, and 
that phosphoric acid is hard to reduce while sulphuric acid is re- 
duced with comparative ease. 

Exercises. — 1. What are the valences of the non-metals in: 
H 2 S 2 7j H 2 Cr 2 7 , KMn0 4 , KH 2 P0 2 , H 3 N0 4 , NaH 2 P0 3 , Na 2 P0 3 ? 
Name these substances. 

2. Is it oxidation or reduction, or neither, when we make, (a) 
N 2 4 from HN0 3 , (6) S0 2 from H 2 S0 3 , (c) HP0 3 from H 3 P0 3 , (d) 
H 2 S 2 7 from H 2 S0 4 , (e) Na 2 S0 4 from NaHS0 3 ? 

3. Why would a mixture of potassium dichromate and hydro- 
chloric acid (p. 270) be less suitable than nitric acid, as an oxidizing 
agent for making phosphoric acid from red phosphorus? 

4. Why is not the tertiary phosphate of sodium (p. 371) decom- 
posed by heating? What tertiary phosphates would be decom- 
posed by this means? 

5. Formulate the hydrolyses of the secondary and tertiary 
sodium orthophosphates as was done for sodium sulphide (p. 271). 

6. How should you prepare Ca 2 P 2 7 and Ca(P0 3 ) 2 , both in- 
soluble? 

7. What product should you confidently expect to find after 
heating (p. 371), (a) sodium phosphite, Na 2 HP0 3 , (b) potassium 
hypophosphite? Make the equations. 

8. Compare the elements chlorine and phosphorus after the 
manner of the comparisons on p. 373. 



CHAPTER XXVIII 
CARBON AND THE OXIDES OF CARBON 

The majority of the substances composing, or produced by, 
living organisms, such as starch, fat, and sugar, are compounds of 
carbon. Hence the chemistry of these compounds is known as 
organic chemistry. It was at first supposed that the artificial 
production of such compounds, e.g., without the intervention of life, 
was impossible. But many natural organic products have now 
been made from simpler ones or from the elements, and the prepa- 
ration of the others is delayed only in consequence of difficulties 
caused by their instability and complexity. On the other hand, 
hundreds of compounds unknown to animal or vegetable life, in- 
cluding many valuable drugs and dyes, have now been added to 
the catalogue of chemical compounds. More than 200,000 differ- 
ent compounds containing carbon are known, and thousands are 
added every year. 

The elements entering into carbon compounds are chiefly hydro- 
gen and oxygen. After these, nitrogen, phosphorus, the halogens 
and sulphur may be named. 

Carbon C 

Occurrence. — Large quantities of carbon are found in the free 
condition in nature. The diamond is the purest natural carbon. 
Graphite, or plumbago, which is the next purest, is found in 
limited amounts, and is a valuable mineral. Coal occurs in 
numerous forms containing greatly varying proportions of free 
carbon. Small quantities of the free element have been found in 
meteorites. 

In combination, carbon is found in marsh-gas, or methane CH4, 
which is the chief component of natural gas. The numerous com- 
pounds found in plants and animals have already been mentioned. 
The mineral oils consist almost entirely of mixtures of various 
compounds of carbon and hydrogen. Whole geological formations 

375 



376 



COLLEGE CHEMISTRY 



are composed of carbonates of common metals, particularly calcium 
carbonate or limestone. 

Allotropic Forms of Carbon. — The allotropic (p. 222) forms 
of carbon differ very strikingly in their physical properties. The 
diamond is transparent, crystalline, and very hard (sp. gr. 3.5). 
Graphite is black, lustrous, and very soft (sp. gr. 2.3). Amorphous 
carbon is very variable. Thus lampblack (see p. 398) is a fine 
powder of nearly pure carbon, charcoal (see p. 408) shows the 
structure of the wood, and coal (see p. 409) contains compounds of 
carbon as well as the free element. These amorphous forms can 
best be discussed after the materials from which they are formed 
have been considered. 

That all the forms are composed of the same element is shown 
by the fact that they all burn in oxygen to give carbon dioxide. 
Then, too, when heated strongly in absence of air, diamond and 
the amorphous forms all turn into graphite. They contain differ- 
ent amounts of chemical energy, however. Thus, when 1 g. of 
each is burned, diamond gives 7805 cal., graphite 7850 and sugar 
charcoal (p. 286) 8040. The tendency of most carbon compounds, 
when heated, to char, giving free carbon, is used as a test. 



The Diamond. — Diamonds are found chiefly in Brazil, and 
South Africa. They are separated by weathering the rock, which 

then crumbles, and by washing the 
debris with water. They are covered 
with a crust which entirely obscures 
their luster, and possess natural crys- 
talline forms belonging to the regu- 
lar system, such as the octahedron 
(p. 83). It should be noted that 
this natural form bears no relation 
whatever to the pseudo-crystalline 
shape which is conferred upon the stone by the diamond-cutter. 
The natural stone is "cut," by grinding new faces. Thus, a 
"brilliant" possesses one rather large, flat face, which forms the 
base of a many sided pyramid (Fig. 99, showing two views). This 
form is given to the stone, in order that the maximum reflection 
of light from its interior may be produced. The diamond is harder 




Fig. 99. 



CARBON AND THE OXIDES OF CARBON 



377 



(Appendix II) than any other variety of matter, so that it can be 
scratched or polished only by rubbing with diamond powder. It 
is the densest form of carbon (sp. gr. 3.5). The colorless stones, 
and occasional specimens with special tints (like the blue, Hope 
diamond) are the most valuable. The black (" carbonado") and 
discolored specimens are used for grinding and glass cutting. 
Mounted round the edge of a tube, they are used for drilling rock, 
so that a cylindrical specimen of the whole of the strata can be 
secured for examination. The forms of carbon are insoluble in all 
liquids at room temperature. Molten iron (q.v.) dissolves five or 
six per cent, part of which goes into combination; but usually only 
graphite is found in the cooled product. Moissan (1887), however, 
succeeded in preparing microscopic fragments of diamonds in this 
way. The diamond is a nonconductor of electricity. 

Diamonds are sold by the new international carat, 200 mgms. 
(old carat, 4 grains = 205 mgms.), and the value increases with 
the size. Thus, a first quality, cut stone of 1 carat is worth about 
$270, one of 2 carats about $340 per carat. The largest diamond 
known, the Cullinan (1905), weighed 3032 (old) carats (621 g. or 
1.37 lbs). It was presented by the Transvaal government to 
King Edward VII, and was cut into stones of 516.5 and 309 
carats and many smaller ones. Other large stones are the Jubilee 
(239 carats), and the Kohinoor (106 carats). 



Graphite. — Graphite (Gk., I write) is found in Cumberland, 
Siberia, Canada, and Ceylon. It is composed of glittering, slip- 
pery, crystalline scales (hexa- 
gonal system). In utter con- 
trast to the diamond, the min- 
eral is extremely soft, has a 
smaller specific gravity (2.3), 
and conducts electricity. It is 
made artificially by an electro- 
thermal process (cf. p. 363). A 
powerful alternating current is 
passed through a mass of granular anthracite, mixed with pitch and 
a little sand (Acheson's process). The mixture (3 tons) is piled 
between the electrodes (Fig. 100) and, on account of its high resist- 
ance, becomes strongly heated. The change occupies 24-30 hours. 




Fig. 100. 



378 COLLEGE CHEMISTRY 



Graphite is now used exclusively for making the anodes in the 
electrolytic manufacture of chlorine and in related processes. 
Mixed with fine clay it forms the "lead" of lead pencils, first used 
in the sixteenth century.* Mixed with clay it is used also for 
making crucibles, which withstand high temperatures and serve 
for melting and casting steel and high melting alloys. As "black- 
lead" it forms stove polish, the layer of fine scales protecting the 
iron against rusting. It is employed as a lubricant in cases where 
oil would be decomposed by the heat and where wooden surfaces 
are in contact. 

Chemical Properties of Carbon, — The most common uses 
of carbon depend upon its great tendency to unite with oxygen, 
forming carbon dioxide C0 2 . Under some circumstances carbon 
monoxide CO (see below) is produced. Aside from the direct 
employment of this action for the sake of the heat which is liber- 
ated, it is used also in the reduction of ores of iron, copper, zinc, and 
many other metals. When, for example, finely powdered cupric 
oxide and carbon are heated, copper is obtained. The gas given 
off is either carbon dioxide, or a mixture of this with carbon mon- 
oxide, according to the proportion of carbon used: 

CuO + C-> Cu + CO, 
2CuO + C->2Cu + C0 2 . 

At the high temperatures produced in the electric furnace, carbon 
unites with many metals and some non-metals. Compounds 
formed in this way are known as carbides, such as aluminium 
carbide AI4C3, calcium carbide CaC 2 , and carborundum CSi (see 
below) . 

The union with hydrogen is ordinarily too slow to be observed. 
But when the carbon is mixed with pulverized nickel (contact 
agent), and hydrogen is passed over the mixture at 250°, methane 
CH4 is formed (99 per cent). The action is reversible and exo- 
thermal, and is therefore, at higher temperature, less complete 
(cf. p. 189), at 850° reaching only 1.5 per cent. On the other hand, 
an electric arc, between carbon poles in an atmosphere of hydrogen, 
gives traces of acetylene C2H2, this action being endothermal. The 

* Priestley was the first to suggest the use of caoutchouc (raw rubber) as 
an eraser. 






CARBON AND THE OXIDES OF CARBON 379 

other compounds of carbon and hydrogen are all obtained by in- 
direct reactions. 

Carbon Bisulphide CS 2 > — This compound is made by direct 
union of sulphur vapor and glowing charcoal. An electric furnace 
like that in Fig. 98 (p. 363) is employed. The substance comes off 
as a vapor and is condensed. 

Carbon disulphide is a colorless, highly refracting liquid (b.-p. 
46°). Traces of other compounds give the commercial article a 
disagreeable smell. It burns in air, forming carbon dioxide and 
sulphur dioxide. It is an important solvent for sulphur and 
caoutchouc (rubber), and dissolves iodine and phosphorus freely. 
Large quantities are employed also in the destruction of prairie 
dogs and ants and for freeing grain elevators of rats and mice. 

Carbon Tetrachloride CCl±. — This compound is manu- 
factured by leading dry chlorine into carbon disulphide containing 
a little iodine (contact agent) in solution: 

cs 2 + 3Ci 2 ^cci 4 + s 2 ci 2 . 

The carbon tetrachloride (b.-p. 77°) is first distilled off, and the 
sulphur monochloride (b.-p. 136°) is purified for use in vulcanizing 
rubber. 

Carbon tetrachloride is a colorless liquid which dissolves fats, 
tars, and many other organic compounds. It is used to take the 
oil or grease out of wool, linen, oil-bearing seeds and bones. It has 
the advantage over gasoline (petrol) and benzine (see p. 391), 
which can be used for similar purposes, that it is non-inflammable. 
"Carbona," used for removing stains from clothing, gloves and 
shoes, is benzine to which sufficient carbon tetrachloride has been 
added to render the mixture noninflammable. "Pyrene" fire 
extinguishers contain, mainly, carbon tetrachloride. The tem- 
perature of the burning material is lowered, because heat is con- 
sumed in vaporizing the liquid, and, at the same time, the vapor 
displaces the air and stops the combustion. 

Calcium Carbide CaC% and Carborundum SiC. — Calcium 
carbide is manufactured in an electric furnace, by the interaction 
of finely pulverized limestone or quicklime with coke: 

CaO + 3C->CaC 2 + CO. 



380 



COLLEGE CHEMISTRY 



The operation is a continuous one, the materials being thrown into 
the left side of the drum (Fig. 101, diagrammatic), and the product 
removed on the right. The carbon poles are fixed. The arc having 
been established, the drum is rotated slowly as the carbide accu- 
mulates. The current enters by one carbon, passes through the 
carbide, and leaves by the other. The high resistance of the 
partially transformed material causes the production of the heat. 
When the action in one layer approaches completion, the resistance 
falls, the current increases, and an armature round which the wire 
passes (not shown in Fig. 101) comes into operation and turns the 
drum. In this way the carbide just formed is continuously moved 

away from the carbons, and 
new material, introduced on 
the left, falls into the path 
of the current. The iron 
plates which form the cir- 
cumference of the drum are 
added on the left and re- 
moved on the right, where 
also the carbide is broken 
out with a chisel. The drum 
revolves once in about three 
days. The product is used 
for making acetylene (q.v.). 

Carborundum, or carbide 
of silicon SiC, of which 
hundreds of tons are manu- 
factured annually at Niagara Falls (Acheson's process), is made 
in an electric furnace of the type shown in Fig. 100 (p. 377). A 
mixture of coke and sand (silicon dioxide Si0 2 ) with some saw- 
dust is piled between the terminals, with a core of granular carbon 
to carry most of the current. The resistance produces a high 
temperature (1950°), and carbon replaces the oxygen: 

3C + Si0 2 ->SiC + 2CO. 

The carborundum remains, often in beautifully crystalline form. 
It is exceedingly hard (Appendix II), and after pulverization and 
mixing with a filler, is moulded into grinding wheels. 




Fig. 101. 



CARBON AND THE OXIDES OF CARBON 381 

Carbon Dioxide and Carbonic Acid 

Occurrence. — Carbon dioxide is present in the atmosphere, 
and issues from the ground in large quantities in certain neighbor- 
hoods, as, for example, in the so-called Valley of Death in Java, 
and in the Grotta del Cane near Naples. Effervescent mineral 
waters, such as those of Vichy and of the Geyser Spring at Sara- 
toga, contain it in solution, and their effervescence is caused by the 
escape of the gas when the pressure is reduced. 

Modes of Formation. — 1. Carbon dioxide is produced by 
combustion of carbon with an excess of oxygen: C + 2 — > C0 2 . 
The combustion of all compounds of carbon, as well as the slow 
oxidation in the tissues of plants and animals, yield the same 
product. The product from burning carbon is naturally mixed 
with at least four times its volume of atmospheric nitrogen. To 
secure carbon dioxide for commercial purposes from this source, 
the gas is led under pressure into a solution of potassium carbonate, 
which absorbs the carbon dioxide: 

C0 2 (gas) <=> C0 2 (dslvd) + H 2 <± H 2 C0 3 + K 2 C0 3 <=± 2KHC0 3 . 

When the pressure is reduced by a pump, all the actions are 
reversed, and the gas escapes in pure form. The same solution, 
with occasional purification, can be used an indefinite number of 
times. 

2. It was Joseph Black (1757) who first recognized the gas as a 
distinct substance. He observed its formation when marble or 
magnesium carbonate was heated : 

CaC0 3 ^CaO + C0 2 , 

and named the gas "fixed air" from the fact that it was contained 
in these solids. The above action had been used for centuries 
in making quicklime (calcium oxide). All common carbonates, 
excepting the normal carbonates of potassium and sodium, de- 
compose, leaving the oxide of the metal or the metal (p. 60). 

3. Black found that the gas was also produced when acids acted 
upon carbonates, and this is the method employed in the labora- 
tory: 

CaC0 3 (solid) <=? CaC0 3 (dslvd) t=5 Ca++ + C0 3 = 1 ^_ rn ^_ r 

2HC1 (dslvd) *=* 2C1- + 2H+ J ^ H2C ° 3 ** Hz ° + C0 " 



382 COLLEGE CHEMISTRY 

Since the carbonic acid is very slightly ionized, the action is like 
that of acids on sulphites (p. 275). Since, however, the carbonate 
of calcium (marble) is very slightly soluble, so that an additional 
equilibrium controls its solution, the action is like that of acids on 
ferrous sulphide (p. 272) . The apparatus shown in Fig. 24 is used. 
Carbon dioxide is formed in decay (p. 36) and, as Black likewise 
discovered, in fermentation (q.v.). 

Physical Properties. — Carbon dioxide is a colorless, odorless 
gas. It is heavier than air. The G.M.V. weighs 44.26 g. Its 
critical temperature is 31.35° (p. 79). The solid melts at —56°, 
having a vapor pressure of 5.3 atmospheres. The solid has a vapor 
pressure of 1 atmos. at —79°. The sp. gr. of the liquid at 0° is 
0.95. At 0° its vapor tension is 35.4 atmospheres and at 20°, 
59 atmospheres. It must be preserved, therefore, in very strong 
cylinders of mild steel. Large quantities of it, often collected from 
fermentation vats, are sold in such cylinders, and used in operating 
beer-pumps and in making aerated waters. When the liquid is 
allowed to flow out into an open vessel or, still better, into a cloth 
bag (non-conductor of heat), it cools itself by its own evaporation 
and forms a white, snowlike mass. Solid carbon dioxide evapo- 
rates at —79°, without melting, since at that temperature it 
exercises 1 atmosphere pressure, and the heat from the surround- 
ings is used as heat of vaporization instead of being employed in 
raising the temperature to the melting-point ( — 56°). 

The solid is used in the laboratory as a cooling agent, being often 
mixed with ether to give closer contact with the vessel. Mercury 
(m.-p. —40°) is easily frozen by the mixture. 

Carbon dioxide gas (760 mm. and 15°) dissolves in its own 
volume of water. Up to four or five atmospheres Henry's law 
(p. 128) describes its solubility accurately. An aqueous solution, 
under a pressure of 3-4 atmospheres, is familiarly known as soda 
water, or carbonated water. 

Chemical Properties. — Carbon dioxide is a stable compound. 
At 2000°, the dissociation reaches 1.8 per cent, or about the same 
as that of water: 2C0 2 <=* 2CO + 2 . 

The more active metals, like magnesium, burn brilliantly when 
ignited in a hollow lump of solid carbon dioxide, producing the 



CARBON AND THE OXIDES OF CARBON 383 

oxide and free carbon. Less active metals, such as zinc and iron, 
when heated in a stream of the gas, give an oxide of the metal and 
carbon monoxide (q.v.). 

Carbon dioxide unites directly with many oxides, particularly 
those of the more active metals, such as the oxides of potassium, 
sodium, calcium, etc. Hence the decomposition of calcium car- 
bonate by heating (p. 381) is a reversible action. 

Carbon dioxide, when dissolved in water, forms an unstable acid: 

H H-0 

H 2 + C0 2 <=± H 2 C0 3 , or ,0 + Cf -» ;C = 0. 

w ^o H-cr 

The name carbonic acid is frequently, though improperly, given to 
the anhydride C0 2 , which has no acid properties. 

Chemical Properties of Carbonic Acid H 2 C0 3 . — The solu- 
tion of carbon dioxide in water exhibits the properties of a weak 
acid. It conducts electricity, although not well. It turns litmus 
red. The ionization takes place chiefly according to the equation: 

H 2 C0 3 ^H+ + HC0 3 - 

Carbonates and Bicarbonates, — When excess of an aqueous 
solution of carbonic acid is mixed with a solution of a base like 
sodium hydroxide, or, as the operation is more usually performed, 
when carbon dioxide is passed into a solution of the alkali, until the 
liquid is saturated, water is formed and the acid carbonate (bi- 
carbonate) of sodium remains dissolved: 

H 2 C0 3 + NaOH fc; H 2 + NaHC0 3 , or H+ + OET->H 2 0. 

Although the bicarbonate is technically an acid salt, its solution is 
neutral, on account of the exceedingly slight dissociation of the 
HC0 3 ~ ion. By addition of an equivalent of sodium hydroxide, 
the normal carbonate is obtained: 

NaOH+NaHC0 3 ^±H 2 0+Na2C0 3 , or OH-+HCOr<=±H 2 0+C0 3 =. 

This solution, like that of all salts of a strong base and a feeble acid 
(c/. p. 271), is alkaline in reaction. This is because the tendency to 
form the very slightly ionized HC0 3 ~ makes the foregoing ionic 
action noticeably reversible (cf. pp. 271, 370). 



384 COLLEGE CHEMISTRY 

The normal carbonates, with the exception of those of potassium, 
sodium, and ammonium, are insoluble in water, and may be ob- 
tained by precipitation when the proper ions are employed. For 
example : 

BaCl2+Na>C08?=*BaC04 +2NaCl, or Ba+++C0 3 =^±BaC0 3 l. 

The aqueous solution of carbon dioxide interacts with solutions 
of barium and calcium hydroxides in a similar manner: 

Ca(OH) 2 + H 2 C0 3 U CaC0 3 j + 2H 2 0. 

These precipitations are used as tests for carbon dioxide. 

Excess of carbon dioxide converts calcium carbonate into the 
more soluble bicarbonate, and hence considerable quantities of 
"lime" (hardness, q.v.) are frequently held in solution by natural 
waters which contain carbon dioxide in solution : 

H 2 C0 3 + CaC0 3 <=> Ca(HC0 3 ) 2 . 

In the same fashion, the carbonates of iron (FeC0 3 ), magnesium, 
and zinc are somewhat soluble in water containing free carbonic 
acid. In fact, the solution, transportation, and deposition of all 
these carbonates take place in nature on a large scale by the alter- 
nate progress and reversal of this action. 

Uses of Carbon Dioxide. — The use of the gas for impregnat- 
ing aerated waters has been mentioned. The gas is used in im- 
mense quantities in the manufacture of sodium bicarbonate 
NaHC0 3 (baking soda), of sodium carbonate Na 2 CO 3 ,10H 2 O 
(washing soda), and of white lead, a basic carbonate of lead 
Pb 3 (OH) 2 (C0 3 ) 2 . _ 

Since carbon dioxide is already fully oxidized, it does not burn, 
and since it is very stable, ordinary combustibles will not burn in 
it. A small percentage of it will destroy the power of air to sup- 
port combustion. For this reason, portable fire extinguishers 
contain a dilute solution of sodium bicarbonate, and a bottle of 
sulphuric acid. When the tank is inverted, the acid flows into the 
solution: 

2NaHC0 3 + H 2 S0 4 <=± Na 2 S0 4 + 2H 2 C0 3 <=± 2H 2 + 2C0 2 . 

The liquid is saturated with the gas and the excess, rising to the 
top, by its pressure forces the solution out through the nozzle. The 



CARBON AND THE OXIDES OF CARBON 385 

liquid is more effective than an equal amount of water, because 
the carbon dioxide it carries mixes with the surrounding air. 

The most wonderful chemical change which carbon dioxide 
undergoes is perhaps the most useful to mankind, and at the same 
time the one least understood. This is the action by which plants 
use it as food (see last section of this chapter). 

Carbon Monoxide CO 

Preparation. — In the laboratory, carbon monoxide is ob- 
tained by heating oxalic acid, a solid, white, crystalline substance, 
in a flask with concentrated sulphuric acid. The latter is here 
employed simply as a dehydrating agent (p. 286), so that it need 
not be included in the equation: 

H 2 C 2 4 -> C0 2 + CO + H 2 0. 

To obtain pure carbon monoxide from this mixture, it is necessary 
to remove the carbon dioxide, by passing the gas through a solu- 
tion of potassium hydroxide contained in a wash bottle. By using 
formic acid, or sodium formate, with sulphuric acid, the presence 
of the carbon dioxide is avoided: 

HCH0 2 ->CO + H 2 0. 

We commonly observe the blue flame of burning carbon mon- 
oxide playing on the surface of a coal fire. The gas is produced by 
the passage of the carbon dioxide, which is first formed, through 
the upper layers of heated coal: 

C0 2 + C->2CO. 

A similar reduction of carbon dioxide is produced when the gas is 
led over a metal, such as zinc, and heat is applied: 

C0 2 + Zn-^ZnO + CO. 

Producer Gas and Water Gas. — When coke and air are 
used in the reaction mentioned above, the mixture of carbon 
monoxide (about 33 per cent) and nitrogen (about 66 per cent) 
obtained is called producer gas. It is combustible and is used in 
factories for heating and to drive gas engines for power. 



386 COLLEGE CHEMISTRY 

When steam is driven through white hot coke or anthracite, a 
mixture of hydrogen and carbon monoxide, known as water gas, 
is produced* 

C + H 2 -* CO + H 2 - 28,300 cal. 

The coke, piled in a brick-lined, cylindrical structure, is brought 
to vigorous combustion by blowing in air for ten minutes. Then 
steam is substituted for the air. Since the interaction takes place 
with absorption of heat (is endothermal, see equation), in about 
five minutes the coke becomes too cool. Air is then substituted 
for steam, and so on alternately. The gas is collected while the 
steam is turned on, and contains equal volumes of the two gases, 
together with some carbon dioxide (4-7 per cent), nitrogen (4-5 
per cent) and oxygen (1 per cent). The gas is, therefore, almost 
wholly combustible and is used as a source of heat, and for driving 
gas engines to furnish power. It is used also for making illuminat- 
ing gas (q.v.). Since carbon monoxide is more easily liquefied than 
is hydrogen, the latter gas is obtained, for commercial use, by 
passing water gas through a liquefier. 

When both steam and air are driven together over burning coke, 
the latter is able to burn continuously, and a fuel gas which is a 
cross between producer gas and water gas is obtained. 

Fuel gases are used on a large scale in steel works, and other 
factories. They give a uniform and easily regulated heat, they 
leave no ash, and their use involves no labor for stoking. As 
gases, also, they can be used in structures in which coal, as a solid, 
could not be employed. 

Physical Properties, — Carbon monoxide is a colorless gas, 
with a metallic odor and taste (poisonous!). It is very slightly 
soluble in water. Its density is almost the same as that of air. 
When liquefied it boils at -190°. 

Chemical Properties. — All the chemical properties of carbon 
monoxide are referable to the fact that in it the element carbon 
appears to be bivalent: CzzO. The compound is in fact unsatu- 
rated, and combines with oxygen, chlorine, and other substances 
directly. Thus the gas burns in the air, uniting with oxygen to 
form carbon dioxide. Again, iron (q.v.) is manufactured by the 



CARBON AND THE OXIDES OF CARBON 387 

reduction of the oxide of iron by gaseous carbon monoxide in the 
blast furnace : 

Fe 2 3 + 3CO <=* 2Fe + 3C0 2 . 

In sunlight carbon monoxide unites directly with chlorine to form 
carbonyl chloride (phosgene) COCl 2 . It unites with nickel and iron 
to form nickel carbonyl and iron carbonyl (q.v.), respectively. 

The gas is an active poison. When inhaled it unites with the 
haemoglobin of the blood, to the exclusion of the oxygen which 
forms a less stable compound (c/. p. 36). A quantity equivalent 
to about 10 c.c. of the gas per kilo, weight of the animal is sufficient 
to produce death, about one-third of the whole haemoglobin having 
entered permanent ly into combination with carbon monoxide. 
One volume in 800 volumes of air produces death in about thirty 
minutes. This gas is the chief poisonous substance in illuminating 
gas. The poisonous effect of tobacco smoke, when inhaled, is 
partly due to the carbon monoxide produced by incomplete com- 
bustion. Nicotine, although contained in tobacco leaves, is 
unstable, and is decomposed by the heat. Traces of other irritant 
organic compounds, however, are contained in the smoke. 

Carbon Dioxide as Plant Food. — The walls of the cells 
which form the framework of a plant are made of cellulose 
(C 6 Hio0 5 )z. In the cells, especially those in certain parts of the 
plant, granules of starch (C6H 10 O 5 ) y are found. The plant con- 
tains also proteins. These substances contain carbon, hydrogen, 
oxygen, nitrogen, sulphur, and phosphorus, and plant food must 
furnish these elements. Compounds of potassium are also re- 
quired. Hence, in addition to large amounts of water ascending 
through the roots and stem, carrying sufficient quantities of solu- 
ble compounds of the four elements last named, all plants require 
an abundant supply of carbon in absorbable form. Now, this 
comes from atmospheric carbon dioxide, admitted through minute 
openings situated mainly on the surfaces of the leaves. Com- 
parison of the formulae C0 2 and C 6 Hi O 5 shows at once that the 
assimilation of the carbon dioxide of the plant must involve 
reduction. The chlorophyll (green matter) and protoplasm in the 
leaves act upon the carbon dioxide, causing oxygen gas to be 
liberated: 

6C0 2 + 5H 2 + 671,000 cal. -* C 6 H 10 O 5 + 60 2 . 



388 COLLEGE CHEMISTRY 

This action goes on only in sunlight, and if green leaves are placed 
under water saturated with carbon dioxide, oxygen is given off 
and can be collected. 

The enormous amount of energy absorbed in the action, and 
represented in terms of heat in the equation, is furnished by the 
sunlight. It may be added that plants, like animals, also use 
some oxygen and produce some carbon dioxide, but this process is 
entirely overborne in daylight, and is noticeable only in the dark. 

The energy that does the world's work comes mainly from two 
sources, namely, water power and the combustion of wood or coal 
(which is fossil wood). The water comes from vapor, generated 
by the sun's heat, condensed as rain, and ultimately feeding the 
rivers. The source of the energy in wood and coal is now apparent. 
When wood, which is largely cellulose (CeHioOs)*, burns, it gives 
carbon dioxide, water, and heat. In fact, its combustion is 
represented by the above equation, when read backwards. Thus, 
the sunlight, through the machinery of the plant, takes carbon 
dioxide and water, supplies the energy (as light), and gives us 
wood and oxygen. And the wood and oxygen, when burned, give 
us back the original substance, and the equivalent of the original 
energy in the form of heat. Thus, the two sources of energy turn 
out to be the same, namely the sun's rays. 

If, instead of burning the starch of the plant, we consume it as 
food, it goes through several changes instead of one. But the final 
products are the same, namely carbon dioxide and moisture, given 
off through our lungs and skin, and heat and other forms of energy 
such as are developed in animals. Thus, whether we use our 
muscles, a steam engine, or a water turbine to do work, sunlight 
is in each case the ultimate source of the energy employed. 

Exercises. — 1. To which two factors in the interaction of 
calcium carbonate and hydrochloric acid (p. 381) is due the forward 
displacement of all the equilibria? 

2. What will be the excess of pressure inside a bottle of soda 
water when 4 vols, carbon dioxide are dissolved in 1 vol. water? 

3. What volume of liquid carbon dioxide, measured at 0°, will be 
required to give 75 liters of the gas at 0° and 760 mm. pressure? 

4. What are the exact relative weights of equal volumes o£ 
carbon dioxide, carbon monoxide, air, and steam? 



CHAPTER XXIX 
THE HYDROCARBONS. ILLUMINANTS. FLAME 

The compounds of carbon and hydrogen are called the hydro- 
carbons. Hundreds of different hydrocarbons, containing different 
proportions of the two elements are known. The natural oil 
petroleum is a mixture of many substances of this class. 

The Irydrocarbons fall into several distinct series, the chief one 
of which contains methane CH4 as its simplest member. On 
account of the fact that certain members of this set are found in 
paraffin, it is commonly known as the paraffin series. For the 
reason that in this series the carbon has all its four valences em- 
ployed, the members are also called the saturated hydrocarbons. 

Paraffin or Saturated Series of Hydrocarbons, — The fol- 
lowing is a list of the names, formulae, and boiling-points of seven 
of the simplest hydrocarbons of this series, and of two of the higher 
members of the series: 



Methane CH< b.-p. 


-164° 


Hexane CeHu b.-p 71° 


Ethane C 2 H 6 " 


- 89.5° 


Heptane C 7 H 16 " 99° 


Propane C 3 H 8 " 


- 37° 


Hexadecane CieH^ " 287.5° 


Butane C4H10 " 


+ 1° 


" m.-p. 18° 


Pentane CsH^ " 


35° 


Pentatricontane C 3 6H 72 " 74.7° 



After the first four, the names are based on the Greek numerals 
representing the number of carbon atoms in the molecule. Hep- 
tane is followed by octane C 8 Hi 8 , nonane C9H20, decane C10H22, 
etc. On examining the formulae, we perceive that, in each, the 
number of hydrogen atoms is equal to twice the number of carbon 
atoms plus two. The general formula is therefore C„H 2 n + 2- 
The series illustrates strikingly the law of combining weights (p. 
42). We note, also, that the first four are gases at room tempera- 
ture. The members from pentane to pentadecane C15H32 are 
liquids, and from hexadecane onwards they are solids. 

In these compounds the carbon is quadrivalent, and each sub- 

389 



390 COLLEGE CHEMISTKY 

stance is related to the preceding one by containing the additional 
units CH 2 . The graphic formulae of the first three members 
illustrate these two facts: 

H H H H H H 

I II III 

H-C-H H-C-C-H H-C-C-C-H 

I II III 

H H H H H H 

These hydrocarbons are extremely indifferent in their chemical 
behavior. They have none of the properties of acids, bases, or 
salts. The halogens, notably chlorine and bromine, however, in- 
teract with them (see below). When burned they all produce 
carbon dioxide and water. 

Petroleum. — Petroleum is a thick, often greenish-brown 
colored oil. When borings reach the oil-bearing strata, the oil, 
hitherto held beneath impervious strata, and often under hydro- 
static pressure of water underneath or around it, either gushes up 
or is pumped to the surface. Wells are in operation in Caucasia, 
Gallicia, India, Japan, and in Ontario, Ohio, Pennsylvania, Cali- 
fornia and elsewhere in North America. The world's production 
in 1912 was 350 million barrels (42 gal. each), of which nearly 220 
millions were produced in the United States. 

The oil is a complex mixture, and is partially separated by dis- 
tillation (p. 93) into products which are still mixtures, but are 
suited to special purposes. The components of lower boiling- 
point come off first and the temperature rises steadily as these 
components are eliminated and those of higher and higher boiling- 
point enter the vapor. As certain temperatures are reached (or 
as the sp. gr. of the distillate attains certain values) the condensed 
liquid is diverted into different vessels, so as to collect together 
the "fractions" of the same kind. This is called fractional dis- 
tillation. 

At some suitable stage, the residual oil is chilled, and a quantity 
of the solid members of the series (C22H46 to C 2 8H 58 ) crystallizes in 
flakes (solid paraffin) and is separated by filtration in presses. 
The final residue is used for lubricants and for fuel. The fractions 
are still mixtures, but contain mainly compounds lying close to- 
gether in the series. Some of the products are as follows: 






THE HYDROCARBONS 



391 



Name. 



Petroleum ether 
Gasoline . . . . 
Naphtha .... 
Benzine . . . . 
Kerosene . . . 



Components. 



Pentane, hexane 
Hexane, heptane 
Heptane, octane 
Octane, nonane 
Decane-hexadecane 



B. 


-P. 


40° 


- 70° 


70°- 


- 90° 


80°- 


-120° 


120°- 


-150° 


150°- 


-300° 



Uses. 



Solvent, gas-making 
Solvent, fuel 
Solvent, fuel 
Solvent 
Illuminating-oil 



Petrolatum (vaseline), C22H46 to C23H48, is separated in some 
refineries. Solid paraffin is employed in waterproofing paper, as 
an ingredient in candles, in the laundry, and to cover preserves. 
Kerosene, for oil lamps, is usually the largest fraction. To be 
used safely, it should not give any inflammable vapor below 65° 
(150° F.), which is the legal flash-point in many states. Special 
treatment, such as superheating the vapor under high pressure 
(Rittman's process), is used to increase the proportion of gasoline 
(petrol) for which there is a large and increasing demand. 

Asphalt, a natural mixture of solid hydrocarbons found particu- 
larly in Trinidad, is used in road-making. 



Methane CH±. — Methane is the chief component of natural 
gas (over 90 per cent), which, like the oil, is confined beneath 
impervious strata and is forced out through borings by hydro- 
static pressure. It is found mainly in or near the localities where 
oil is found. It also rises to the surface when the bottoms of 
marshy pools are disturbed (Marsh-gas), and issues from seams in 
coal beds as fire-damp (Ger. Dampf, vapor). In these two cases 
it results from the decomposition of vegetable matter in absence 
of air. Its formation by direct union of carbon and hydrogen has 
already been discussed (p. 378). 

Methane may be made from inorganic materials by the action of 
water upon aluminium carbide, prepared by the interaction of 
aluminium oxide and carbon in the electric furnace (c/. pp. 378, 
377): 

AI4C3 + 12H 2 -> 4A1(0H) 3 + 3CH4. 

In the laboratory the gas is commonly obtained by the distillation 
of a dry mixture of sodium acetate and sodium hydroxide: 

NaC0 2 CH 3 + NaOH -> Na 2 C0 3 + CH4. 



392 COLLEGE CHEMISTRY 



As regards chemical properties, methane, like other saturated 
hydrocarbons (p. 390), is very indifferent. When a mixture of 
methane and chlorine is exposed to sunlight several changes occur 
in succession (cf. p. 162) : 

CH4 + Cl 2 -* CH 3 C1 + HC1, CH 3 C1 + Cl 2 -» CH 2 C1 2 + HC1, 
CH 2 C1 2 + Cl 2 -> CHCI3 + HC1, CHCI3 + Cl 2 -+ CC14 + HC1. 

This kind of interaction with the halogens is characteristic of com- 
pounds of hydrogen and carbon. It takes place slowly, and is 
therefore entirely different from ionic chemical change. It con- 
sists in a progressive substitution of chlorine for hydrogen, unit by 
unit. Chloroform CHCI3 and carbon tetrachloride CCI4 (p. 379) 
are familiar substances. The iodine derivative, iodoform CHI 3 is 
used in surgical dressing. These substances are not salts, and are 
not ionized in solution. They are very slowly hydrolyzed by 
water — carbon tetrachloride, for example, giving carbonic acid 
and hydrochloric acid. 

Methane and the other saturated hydrocarbons are decomposed 
by strong heating (see cracking, below). 

Unsaturated Hydrocarbons. — In addition to the saturated 
series of hydrocarbons, several other series are known in which 
smaller proportions of hydrogen are present. Thus, ethylene 
C 2 Hi, to which illuminating gas largely owes the luminosity of its 
flame, belongs to a series C n H 2ra , all the members of which contain 
two atoms of hydrogen less than the corresponding compounds of 
the first series. Again, acetylene C 2 H 2 is the first member of a series 
C n H 2n _ 2 , and benzene CeH 6 begins a series C n H 2n _ 6 , of which toluene 
C 7 H 8 (p. 349) is the second member.* These are all unsaturated 
because the full valence of the carbon is not in use, and these 
compounds, therefore, unite more or less readily with hydrogen, 
chlorine, bromine, and concentrated sulphuric acid. The hydro- 
carbons of all the series are mutually soluble, but none of them 
dissolve in water. 

Members of the ethylene and acetylene series are found in 

* Isoprene C 6 H 8 , a member of the unsaturated series C71H271-2, when heated 
in presence of sodium (or some other contact agent), changes into caoutchouc 
(C 6 H 8 ) X or raw rubber. No method of preparing synthetic rubber has yet been 
used commercially. 



. 



THE HYDROCARBONS 393 

petroleum, and are formed also to some extent by decomposition 
during the distillation. As oil containing them acquires dark- 
colored products by chemical change, the oils are always refined 
before being sold. They are agitated with concentrated sulphuric 
acid, which unites with the unsaturated substances and, being 
insoluble in the oil, collects in a layer below it. The oil is finally 
washed free from the acid with dilute alkali and with water. 

Ethylene C 2 ff 4 . — Ethylene is the first member of the second 
series of hydrocarbons. It corresponds to ethane C 2 H 6 , but con- 
tains in each molecule two hydrogen units less than does this 
substance. 

Ethylene is made by heating common alcohol (ethyl alcohol) 
with concentrated sulphuric acid: 

C 2 H 5 OH->H 2 + C 2 H4T. 

A comparison of the graphic formulae of the alcohol and ethylene 
shows that this loss of water leaves the carbon partly unsaturated: 

H H H H H H 

II II II 

H-C-C-O-H -C-C- or H-C = C-H 

II II 

H H H H 

The elements of water may also be removed by allowing the alcohol 
to fall drop by drop onto heated phosphoric anhydride. 

Ethylene is formed, along with acetylene and other substances, 
when any saturated hydrocarbon is heated strongly. Even 
methane gives it: 

2CH4 — > C 2 H4 -f- 2H 2 . 

Ethylene is a gas. It burns in the air with a flame which, on 
account of the great separation of free carbon which takes place 
temporarily during the combustion (cf. Flame), is highly luminous. 
It will be seen that, in the formula, but three of the valences of each 
carbon unit are occupied: the substance is unsaturated. Hence, 
when ethylene is passed through liquid bromine it is rapidly 
absorbed, and the bromine seems to increase in volume and finally 
loses all its color, being converted into a transparent liquid having 
the composition C 2 H4Br 2 , ethylene bromide. 



394 COLLEGE CHEMISTRY 

Acetylene. — This substance, likewise a gas, is the first member 
of still another unsaturated series. Its formula C 2 H 2 shows that 
its molecule lacks four of the hydrogen units necessary to the com- 
plete saturation which we find in ethane. Graphically its structure 
is usually represented thus : H-CEC-H. This gas is 
formed in small quantities by direct union of carbon and hydrogen 
in the electric arc (p. 378). This is because the reaction is en- 
dothermal (p. 189). For the same reason, it is also produced when 
ethylene is passed through a heated tube : C2H4 — * C 2 H 2 + H 2 
(c/. Flame). 

When calcium carbide (p. 379) is thrown into water it is hy- 
drolyzed. Violent effervescence occurs, the calcium carbide is 
disintegrated, a precipitate of calcium hydroxide is formed, and 
acetylene passes off as a gas: 

CaC 2 + 2H 2 -> Ca(OH) 2 + C 2 H 2 . 

This action is like that of water on calcium phosphide (p. 365), 
magnesium nitride (p. 339), and aluminium carbide (p. 391). 

Acetylene burns with a flame which is still more luminous than 
that of ethylene. On account of the large amount of heat absorbed 
when it is formed: 2C + H 2 — > C 2 H 2 — 53,200 cal., an equal 
amount is liberated when it decomposes. If the gas is compressed 
in tanks, it is therefore apt to explode from any shock. It is 
frequently made in generators, as needed, by the foregoing action, 
and is used for lighting on automobiles and in regions remote from 
a public supply of illuminating gas. The acetylene tanks, which 
are also in use, contain acetylene dissolved under high pressure 
in acetone, a form in which it can be handled safely. 

When acetylene C 2 H 2 is burned, we obtain from 2 X 12 -f 2 = 
26 g. not only the heat due to the combustion of the carbon (2 X 
12 X 8040 cal., p. 376), and of the hydrogen (2 X 28,800 cal.), but 
also the heat due to the decomposition of the gas (53,200 cal.). 
The temperature of the flame is therefore extraordinarily high. 
The oxyacetylene flame, produced by means of a suitable burner 
(Fig. 32, p. 58), is now used, under the name of the acetylene 
torch, for cutting metals. The gases are contained in portable 
tanks. Such a flame will melt its way through a 6-inch shaft or a 
steel plate several feet wide in less than a minute, cutting the object 
in two. Steel buildings have thus been taken down, and ships 



THE HYDROCARBONS 395 

(like the Maine) have been cut up for removal. Other gases, like 
blau gas and oil gas, made by cracking petroleum (see below), are 
now displacing acetylene for this purpose, as they are almost as 
effective, and the flame is more easily controlled. 

Cracking of Hydrocarbons. — All hydrocarbons, when heated 
strongly (air-excluded) decompose or crack. The changes seem 
to be reversible, and the result therefore depends upon the con- 
ditions. Thus, at atmospheric pressure, and especially when the 
oil is mainly present as a liquid, hydrogen is given off and un- 
saturated liquid and gaseous hydrocarbons are produced. Under 
such conditions, ethylene is formed in large amounts. On the 
other hand, when an oil free from gasoline is completely vaporized 
(500°), and is under high pressure, the hydrogen is forced into 
combination with the broken molecules and the saturated con- 
stituents of gasoline are produced in large amount (Rittman's 
process) . 

At a white heat, all the hydrocarbons decompose into hydrogen 
and free carbon. The latter is deposited in a dense form called 
gas-carbon, which is more or less crystalline (like graphite) and 
used in making carbon rods for arc lights and electric furnaces, 
and carbon plates for batteries, and for the electrodes employed 
in electrolysis. The carbon is ground up, moistened with petro- 
leum residues, subjected to hydraulic pressure and finally heated 
strongly to expel volatile matter. 

Carburetted Water Gas. — As we have seen, water gas is 
essentially H 2 + CO (p. 386), and burns with a pale-blue flame. 
To fit it for use as illuminating gas, unsaturated hydrocarbons, 
which burn with a luminous flame, such as ethylene C2H4 and 
acetylene C 2 H 2 must be added. The gas is sent through a tower 
containing strongly heated brick on which a petroleum oil is 
sprayed. Mixed with the vapor, the gas then passes into the 
"superheater" where, at a higher temperature, the cracking into 
unsaturated hydrocarbons occurs. The gas is then cooled and 
washed to remove condensible hydrocarbons, which would other- 
wise obstruct the service pipes. A typical carburetted water gas 
has the composition: Illuminants 17 per cent; heating gases, 
methane 20 per cent, hydrogen 32 per cent, carbon monoxide 26 



396 COLLEGE CHEMISTRY 

per cent; impurities (nitrogen and carbon dioxide) 5-6 per cent. 
A flame burning 5 cu. ft. per hour gives 25 candle power. 

Blau gas and Oil gas, such as Pintsch gas, contain larger pro- 
portions of illuminants. Thus a good oil gas shows: illuminants 
45 per cent; heating gases, methane 39 per cent, hydrogen 14.5 
per cent; impurities 1.5 per cent; candle power 65. Such gases 
are compressed in tanks and used for illumination on railway 
trains (Coal gas, see p. 410). 

Flame 

Meaning of the Term. — In the combustion of charcoal there 
is hardly any flame, for the light emanates almost entirely from the 
incandescent, massive solid. When two gases are mixed and set on 
fire, a sort of flame passes through the mixture, but this can hardly 
be accounted a flame, in the ordinary sense, either. The rapid 
movement of the flash, and the explosion which accompanies it, 
are in a manner the precise opposite of the quiet combustion which 
is characteristic of flames. 

With illuminating gas the production of its very characteristic 
flame is due to the chemical union of a stream of one kind of gas 
in an atmosphere of another. The flame is made up 
of the heated matter where the two gases meet. In 
the case of a burning candle (Fig. 102), one of the 
bodies appears to be a solid, but a closer scrutiny 
of the phenomenon shows that the solid does not 
burn directly. A combustible gas is manufactured 
continuously by the heat of the combustion and 
rises from the wick. The introduction of a narrow 
tube into the interior of the flame enables us to 
draw off a stream of this gas and to ignite it at a 
remote point. Thus, a flame is a phenomenon produced at the 
surface where two gases meet and undergo combination with the 
evolution of heat and, more or less, light. 

In the chemical point of view, it is a matter of indifference 
whether the gas outside the flame contains oxygen, and the gas 
inside consists of substances ordinarily known as combustibles, or 
whether this order is reversed. In an atmosphere of ordinary 
illuminating gas, the flame must be fed with air. This condition 




Fig. 102. 



FLAME 



397 




Fig. 103. 



is easily realized (Fig. 103). The lamp chimney is closed at the 

top until it has become filled with illuminating gas. This can be 

ignited as it issues from the bottom of the wide, straight tube. 

When the hole in the cover of the lamp chimney 

is then opened, the upward draft causes the flame ^= 

of the burning gas to recede up the tube, and there 

results a flame fed by air and burning in coal gas. 

Luminous Flames, — The flame of hydrogen, 
under ordinary circumstances, is almost invisible, 
nearly all the energy of the combustion being de- 
voted to the production of heat. A part of this, 
however, may be transformed into light by the 
suspension of a suitable solid body, such as a 
platinum wire, in the flame. The holding of a 
piece of quicklime in an oxy hydrogen flame (cf. p. 
58) is a practical illustration of this method of 
securing luminosity. In general, luminosity may 
be produced by the presence of some solid which is heated to 
incandescence. 

In the Welsbach lamp the flame itself is nonluminous and, but 
for the mantle, would be identical with the ordinary Bunsen flame. 
The mantle which hangs in the flame, however, by its incandes- 
cence, furnishes the light. This mantle is composed of a mixture 
of 99 per cent thorium dioxide Th0 2 and one per cent cerium 
dioxide Ce0 2 . These oxides act as a contact agent, hastening 
the combustion and liberation of heat close to their surface, and so 
acquire a temperature higher than the average for the rest of the 
flame. The Welsbach lamp gives four times as much light as does 
the same gas, issuing at the same rate, from an ordinary burner. 

In cases of brilliant combustion, as of magnesium ribbon or 
phosphorus, a solid body is formed whose incandescence accounts 
for the light. The flame of ordinary illuminating gas does not at 
first sight appear to give evidence of the presence of any solid body. 
But if a cold evaporating dish is held in the flame for a moment, a 
thick deposit of finely divided carbon (soot) is formed, and we at 
once realize that the light is due to the glow of these particles in a 
mass of intensely hot gas. Carbon is, indeed, an extremely com- 
bustible substance, and is eventually entirely consumed. But a 



398 



COLLEGE CHEMISTRY 



fresh supply is being generated continuously in the interior of the 
flame, while the oxygen with which it is to unite is outside the 
flame altogether. Thus the carbon particles persist until, drifting 
with the spreading gas, they reach the periphery of the flame. 

On a large scale, oil residues are burned so that the flame strikes 
a revolving, iron vessel cooled with water. The soot or lampblack 
is continuously scraped off as the vessel turns. Lampblack is 
used in making printer's ink, India ink, and black varnish. 



The Bunsen Flame. — In the burner devised by Robert 
Bunsen, a jet of ordinary illuminating gas is projected from a nar- 
row opening into a wider tube (Fig. 104). In this 
tube it becomes mixed with air, entering through 
openings whose dimensions can be altered by means 
of a perforated ring. When the supply of air is 
sufficient, the flame becomes non-luminous. With 
a somewhat different construction, and the use of 
a bellows to force a larger proportion of air into 
the gas, a still hotter flame can be produced. The 
instrument in this case is known as a blast lamp. 

The high temperature of the blast lamp flame 
presents an interesting problem. The same amounts 
of gas and air burn to give the same amounts of 
the same products, whether the air blast is on or 
off. The same amount of heat is produced, and 
the same quantities of the same substances are 
heated. The average temperature throughout the 
flame should therefore be the same. In point of fact, it is the 
same, but the stream of hot gas is moving more rapidly when the 
blast is going. The temperature of a body immersed in the flame 
depends, on the one hand upon the rate at which heat reaches it, 
and upon the other on the rate at which it loses heat by radiation. 
The heat is partly carried by the moving, heated gases (convection), 
and partly transmitted by conduction through the stationary layer 
(p. 331) on the surface of the body. Now, the latter is the 
slower process. Hence a rapid stream of gas, which leaves a 
thinner stationary layer, will diminish the distance the heat has to 
travel by conduction and so convey heat to the body faster than 
could a slow stream of the same temperature. Thus, with a blast 




Fig. 104. 



FLAME 399 

flame, the loss by radiation is the same at the same temperature, 
but heat reaches the body faster and so the temperature of the 
body more nearly approaches that of the flame itself. 

Structure of the Illuminating and the Bunsen Flame. — 

When an exceedingly small luminous flame is examined, the various 
parts of which it consists may easily be made out. In the interior 
there is a dark cone which is composed of illuminating gas and air, 
and in it no combustion is taking place. A match-head may be 
held here for some time without being set on fire. This is there- 
fore not properly a part of the flame. Outside this is a vivid blue 
layer (C, Fig. 105) which is best seen in the lower part 
of the flame, but extends beneath the luminous sheath, 
and covers the dark inner cone completely. Outside 
the blue flame, and covering the greater part of it, is 
the cone-shaped luminous portion (B). Over all is a 
faint mantle of non-luminous flame (A), which be- 
comes visible only when the light from the luminous 
part is purposely obstructed. In the luminous gas- 
flame, therefore, there are four regions, if we count 
the inner cone of gas. The difference between this 
and the non-luminous Bunsen flame is that in the 
latter the luminous region is omitted, and the inner, 
dark cone, the blue sheath, and the outer mantle, 
are the only parts which can be distinguished. 

The Causes of Luminosity and Non-Lumi- 
nosity. — The study of the chemical changes taking 
place in the Bunsen flame, particularly with the 
object of explaining (1) the luminosity of the flame of the pure 
gas, and (2) the non-luminosity of that produced by the same gas 
when it is mixed with air, has been the subject of many elaborate 
investigations. The questions are: (1) Why is carbon liberated 
in the former case, and (2) why is it not liberated in the latter? 
Let us consider these questions in order. 

1. The investigations of Lewes (1892) and others show conclu- 
sively that the free carbon in the luminous zone of the ordinary 
flame is accompanied by free hydrogen, and that both are formed 
by dissociation of the ethylene. This substance, when heated s 



400 COLLEGE CHEMISTRY 

gives acetylene, and the latter then dissociates into carbon and 
hydrogen (p. 394) : 

C2H4 — > II2 -f- C2H2 — > 2C -f* H-2. 

The carbon glows, until, as it drifts outwards, it encounters the 
oxygen of the air and is burned. The first oxygen encountered 
combines more readily with the hydrogen, since it is a gas, than 
with the carbon, which is now in solid particles and therefore burns 
less readily. That carbon glows when heated in the absence of 
oxygen, without being consumed, is a fact familiar in the behavior 
of the incandescent electric lamp, the filament of which is often 
made of carbon. 

The conception that when hydrocarbons burn, they first undergo 
dissociation, and then union with oxygen, is in harmony with what 
we have observed also in the case of the combustion of hydrogen 
sulphide, where the presence of free sulphur and free hydrogen in 
the interior of the flame was demonstrated (p. 268). 

2. The influence of the air admitted to the Bunsen burner, in 
interfering with this dissociation in such a way as to destroy all 
luminosity, is the most difficult point to explain. The 
A effect is frequently attributed to the oxygen which the 

( \ air contains. This view, however, is seriously weak- 

* ened by a consideration of the undoubted fact that 

oxygen is not required. Carbon dioxide and steam 
are equally efficient when introduced instead of air 
(Fig. 106, gas enters at a and C0 2 at 6). Even nitro- 
gen, which cannot possibly be suspected of furnishing 
any oxygen, likewise destroys the luminosity. Lewes 
has shown that 0.5 volumes of oxygen in 1 volume of 
coal gas destroy the luminosity. But 2.30 volumes 
of nitrogen or 2.27 volumes of air accomplish the same 



s v , result. Thus the efficiency of air is not much greater 
than that of nitrogen, in spite of the fact that one- 
fifth of the former is oxygen. 
It is evident that the effect is due, in part at least, to the dilution 
with a cold gas. This is confirmed by the observation that a cold 
platinum dish held in a small luminous flame is similarly destruc- 
tive of the luminosity. If the tube of the Bunsen burner is heated 
so that the mixed gases are considerably raised in temperature 



FLAME 



401 



before reaching the non-luminous flame, the latter becomes lumi- 
nous. It is probable, therefore, that the cold gas lowers the tem- 
perature of the inner flame, and at the same time the dilution 
diminishes the speed with which the free carbon is formed (Lewes) . 
Even if the temperature is not reduced below that at which dis- 
sociation of the ethylene can occur, yet the dilution and cooling, 
together, prevent that sharp dissociation at this particular point 
which is necessary for the production of the great excess of free 
carbon needed to furnish the light. 

Before these investigations were made, a different answer was 
given to the question why the flame of pure illuminating gas con- 
tains free carbon and is luminous. It was said that hydrogen was 
more easily burned than carbon, and therefore the latter was left 
free, to be burned later. It is true that gaseous 
hydrogen burns more easily than solid carbon, e.g., 
charcoal. But in ethylene, both elements are equally 
gaseous and the explanation is faulty. Smithells 
(1892) demonstrated the falsity of this explanation 
by devising a cone-separator (Fig. 107). The air 
and ethylene or other gas are admitted separately, 
and the inner cone of the non-luminous flame rests 
on the inner, narrow tube, while the outer cone is 
at the top. By means of a side tube (not shown) he 
withdrew the inter-conal gas and found that, while 
all of the carbon was burned by the inner cone as far 
as carbon monoxide CO, most of the hydrogen was 
still entirely uncombined. The change in the inner 
cone of the Bunsen flame consists, therefore, mainly 
in the burning of all the hydrocarbons to carbon monoxide, with 
liberation of the hydrogen. In the outer cone, it is practically a 
burning of water gas that is taking place. 




Fig. 107. 



Exercises. — 1. Write a graphic formula for hexane. 

2. Write an equation for the formation of aluminium carbide 
(p. 391). 

3. Make a section showing the shape of the flame produced by 
burning hydrogen gas when the latter issues from a circular 
opening. 



CHAPTER XXX 

THE CARBOHYDRATES AND RELATED SUBSTANCES 

A plant takes carbon dioxide from the air and water from the 
ground and, using the energy of sunlight, converts them into a 
growing framework of cellulose (C 6 Hio0 5 )y and, as we have seen 
(p. 387), into starch (CeHioOs)* which it stores in the cells. 
The cellulose of certain plants furnishes us with cotton, linen, 
jute, and paper. The starch of wheat, oats, maize (corn), and 
potatoes is one of the chief foodstuffs they contain. The plant, 
when dead and buried, changes into coal. The fresh wood, when 
distilled, supplies wood spirit (methyl alcohol) and other useful 
substances, and the residue is the valuable charcoal. Further- 
more, from starch we can readily make sugar, alcohol, and other 
familiar materials. Cellulose, starch, and the sugars (e.g., cane- 
sugar C12H22O11) contain oxygen and hydrogen in the same pro- 
portions in which they are present in water, 2H : 10. They might 
be considered hydrates of carbon, and so they are called the 
carbohydrates. The foregoing brief summary shows that the 
carbohydrates introduce us~to a much greater variety of inter- 
esting organic compounds than does petroleum. 

Cellulose (C e H 10 0s)x and Paper. — The wall of each cell, 
and therefore the whole framework of a plant, is made of cellulose. 
Linen and cotton are pure cellulose. The walls of the cells are 
usually more or less thickened by a substance called lignin, which 
has much the same composition, but different chemical behavior. 
The best paper is made of cotton or linen (rag-paper). Cheaper 
kinds are prepared by cutting wood, such as spruce or pine, into 
chips and treating (" cooking") them with a solution of calcium 
bisulphite Ca(HS0 3 )2. This process decomposes the lignin, and 
converts it into soluble materials. The sulphite liquor is then 
run off, and the pulpy material is washed, beaten with water to 
reduce it to minute shreds, and bleached with very dilute chlorine- 

402 




THE CARBOHYDRATES AND RELATED SUBSTANCES 403 

water. The pure cellulose, now paper-pulp, is suspended in 
water, spread on screens, pressed, and dried. During the process, 
other substances are added. Thus, size (gelatine or rosin and 
alum, see Sizing) prevents the ink from running; pulverized 
calcium sulphate (gypsum), clay, and other white solids ("load- 
ing") give body to the paper and permit the production of a 
smooth surface by rolling (" calendering"). Dyestuffs can be 
added to give special tints. Filter paper is pure cellulose. 

Starch (C 6 H 10 O s )„. — Starch consists of little colorless granules 
of various rounded shapes (Fig. 108), which are easily seen under 
the microscope. These granules are massed in large quantities 
in the ears of wheat and oats, in the 
tubers of potatoes, in the grains of 
maize (corn), and in peas and beans. 
Even in the leaves they can be seen. 

Starch is recognized by the iodine test \(fl||r \!jgr ^^ 

(p. 3), turning deep blue with a trace 
of free iodine. 

The treatment of wheat flour, which FlG " 108 ' 

is three-fourths starch, by washing out the starch through a porous 
cloth with water, has already been described (p. 3). It is made 
from maize in America and from potatoes in Europe, by washing 
the flour on sieves. 

Starch is not soluble in water. If boiled with water, however, 
the granules swell and break and the starch is diffused through 
the water, giving a clear liquid. If too much water is not used, 
the liquid when cold sets as a jelly. While the liquid is hot, much 
of the starch will pass through a filter along with the water. 
Such a liquid is called a colloidal suspension. Suspensions like 
this are constantly met with in using complex organic compounds 
like jellies, glues, soaps, and dyes. Even insoluble inorganic 
materials, like gold, give such suspensions (see p. 416). 

The colloidal suspension of starch turns blue when a solution 
containing free iodine is added to it. It is used in the laundry for 
stiffening white goods. Glucose is manufactured from it. 

Glucose CeHviOe, a Sugar, from Starch. — When starch is 
boiled with water, to which a few drops of an acid (contact agent), 



404 COLLEGE CHEMISTRY 

such as hydrochloric acid, have been added, the liquid, after 
neutralization of the acid, is found to be sweet. One of the sugars, 
glucose CeHi 2 06, can be obtained in crystals after evaporation. 
The crystals form "brewers' glucose" and the syrup produced by 
concentration is corn syrup (if maize is the source of the starch). 
The latter, although less sweet than ordinary sugar, is much less 
expensive and is used in making preserves and cheap candy. 

The molecular weight of starch is unknown, but undoubtedly 
large. The formula (CeHioOs)^ shows the composition. The 
water, in presence of a little acid, decomposes the molecules and 
combines with the material. First, dextrin (used as paste or 
mucilage) is formed and this breaks up into glucose. The action 
is an hydrolysis : 

(C 6 H 10 O 5 ) y + 2/H 2 ->2/C 6 H^0 6 . 

Glucose is known also as dextrose and as grape sugar. The 
crystalline granules in raisins (dried grapes) are composed mainly 
of it. When pure, it is almost colorless. It reduces cupric 
hydroxide, in Fehling's solution (q.v.), to cuprous oxide. 

Corn syrup contains 30-40 per cent of unchanged dextrin, 40-50 
per cent of glucose, and the rest water. 

The Sugars. — The common sugars may be divided into the 
monosaccharides, usually with the formula C 6 Hi 2 6 , and the 
disaccHarides, usually C^H^On. Of these, the following will be 
referred to in what follows: 

Monosaccharides: Glucose (grape sugar or dextrose) C 6 Hi 2 6 , 
Fructose (fruit sugar or levulose) CeH^Oe. 

Disaccharides: Sucrose (cane-sugar, beet-sugar) C12H22OU. 
Maltose (action of malt on starch) C12H22O11. 
Lactose (milk sugar, in animals only) C12H22O11. 

Sucrose, or Cane Sugar. — Plants, such as the sugar-cane 
and beet, besides forming cellulose and starch, produce excep- 
tional amounts of sucrose, or table sugar. The sap of the sugar 
maple contains much of it. 

Cane sugar is extracted by crushing the stalks between rollers, 
and evaporating the expressed liquid (18 per cent sugar) in closed 
pans. A partial vacuum is maintained so that the solution may 



THE CARBOHYDRATES AND RELATED SUBSTANCES 405 

boil at a low temperature (65° to begin with) and none of the 
sugar be decomposed. When the syrup cools, the sucrose ap- 
pears in brown-colored crystals. The mother-liquor is called 
molasses. In the refinery, the sugar is redissolved, the solution is 
poured through a column of charcoal, which takes out the coloring 
matter, and the liquid is once more allowed to deposit crystals. 
Pure cane-sugar has a faint } r ellow tint, and a small amount of 
ultramarine (q.v.) is added to give that whiteness which is pop- 
ularly connected with purity in sugar. 

Sugar beets (16 per cent or more sugar) are sliced and steeped 
in water. The extract contains a gummy material in colloidal 
suspension. This is coagulated and precipitated by adding slaked 
lime (calcium hydroxide) Ca(OH) 2 suspended in water, and boil- 
ing. After separation of the clear liquid, carbon dioxide is passed 
through it to precipitate the excess of lime as carbonate (CaCOs). 
The solution is then decolorized with charcoal and evaporated to 
crystallization. 

As regards properties, sucrose crystallizes in four-sided prisms 
(rock-candy, Fig. 48, p. 83) and melts at 160°. When heated to 
200-210° it partially decomposes, leaving a soluble, brown, mixed 
material, caramel, used in coloring whisky and soups. Sucrose 
does not reduce Fehling's solution. 

When boiled with water containing a trace of acid (contact 
agent), sucrose is hydrolyzed, giving a mixture of the two mono- 
saccharides, glucose and fructose: 

C12H22O11 + H2O — > CeH^Oe + C6H12O6. 

The mixture is called invert sugar, and is found in many sweet 
fruits and in honey. Each sugar interferes with the crystalliza- 
tion of the other, by lowering the freezing point (p. 134), and so 
invert sugar is added in making fondant candy and candy that 
is to be pulled, both of which must remain soft for some time. 
Icing for cakes has to some extent this property, and is made by 
adding acid substances to sugar, such as vinegar, lemon juice, or 
cream of tartar. 

Enzymes. — Yeast, consisting of microscopic cells, belongs to 
a low order of plants. Its use lies in the fact that, while multi- 
plying, it secretes within each cell two very active, soluble sub- 



406 



COLLEGE CHEMISTRY 



stances. These are zymase and sucrase (invertase) which belong 
to a class of organic substances called enzymes. Sucrase means 
an enzyme that splits sugar. Enzymes produce remarkable 
chemical changes by their mere presence (contact actions). 

Alcoholic Fermentation. — When some yeast, which is a 
mass of the living plants, is added to a solution of glucose at 
about 30°, the small amount of zymase gradually decomposes 
the sugar. Bubbles of carbon dioxide soon begin 
to rise, and may be tested (p. 384) with limewater 
(Fig. 109) . At the same time, alcohol (ethyl alco- 
hol C 2 H 5 OH) accumulates in the liquid: 




Fig. 109. 



CeH^Oe -* 2C0 2 1 + 2C 2 H 5 OH. 

Yeast will ferment fructose CeH^Oe with the same 
result, but more slowly, so that, when placed in 
invert sugar, it decomposes the glucose first and 
the fructose afterwards. 

Zymase does not act upon sucrose (table sugar), but sucrase 
hydrolyzes the sucrose in the same way as does a dilute acid, giv- 
ing invert sugar. The latter is then decomposed by the zymase, 
and so cane-sugar in solution is fermented by yeast into alcohol 
and carbon dioxide, just as is glucose, only more slowly. 

In making wine the glucose contained in the grape juice is fer- 
mented by a species of yeast found on the skins of the grapes. 

Commercial Alcohol is not made from sugar, but from the starch 
of potatoes or maize. When barley is allowed to sprout, an 
enzyme, amylase (meaning starch-splitting enzyme) or diastase, 
is formed in the ears. The whole material is dried, and is then 
called malt. When this is mixed with starch and water, the 
amylase hydrolyzes the starch to maltose C12H22O11 (p. 404). 
The latter is then further hydrolyzed by yeast to form glucose 
C 6 Hi20 6 , and this is decomposed by the zymase into alcohol and 
carbon dioxide. 

Whisky (about 50 per cent alcohol) is made by treating the 
starch of rye, maize or barley in the same way, with subsequent 
distillation (see below) to separate the alcohol (whisky). Beer is 
made similarly from various kinds of grain, especially barley, but 
the fermented liquor is not distilled. 



THE CARBOHYDRATES AND RELATED SUBSTANCES 407 

Ethyl Alcohol C 2 H 5 OH. — Common alcohol is related to 
ethane C 2 H 6 , having an hydroxyl group in place of one unit of 
hydrogen. Hence its name, ethyl alcohol. 

Ethyl alcohol boils at 78.3° and so, when the fermented liquor 
is distilled, it is almost pure alcohol that comes off. Commercial 
alcohol contains 95 per cent by volume (in Great Britain, 90 per 
cent). Absolute alcohol is made by adding quicklime, which 
combines with the water, and redistilling the liquid. 

Alcohol mixes with water in all proportions. In dilute, aqueous 
solution it is not ionized, and does not interact with acids, bases, 
or salts. It is, however, easily oxidized to acetic acid. When 
water is absent, it interacts with acids slowly (see p. 413). 

Alcohol is used as a solvent for the resins employed in making 
varnishes for wood and lacquers for metal. 

On account of the high duty on 95 per cent alcohol ($2.11 per 
gallon in the U. S. and 24/6 in Gt. Britain), denatured alcohol 
(methylated spirit), which is free of duty, is employed for indus- 
trial purposes. The alcohol (cost about 22 cents per gal. in the 
U. S. and 1/6 in Britain) is mixed with offensive or poisonous 
materials, which prevent its consumption as a beverage, without 
interfering with other uses. Wood spirit and gasoline are often 
employed. 

Acetic Acid HC0 2 CH 3 . — This is the sour substance in vinegar, 
and has many industrial applications. Vinegar is made by oxi- 
dizing alcohol with atmospheric oxygen, using an enzyme se- 
creted by bacterium aceti (mother of vinegar) as contact agent. 
Dilute alcohol from any source, such as fermented apple juice 
(hard cider), is allowed to trickle over shavings in a barrel. Holes 
admit air, and the shavings are inoculated in advance by wetting 
with vinegar: 

HOC 2 H 5 + 2 -> HC0 2 CH 3 + H 2 0. 

The issuing liquid contains 5-15 per cent of acetic acid, which can 
be purified by fractional distillation to separate the water. It 
boils at 118° and freezes at 16.7°. Although four atoms of hydro- 
gen are contained in its molecule, but one of these is replaceable 
by metals. This fact is recognized in the reaction formula (p. 
95) of the acid, HC 2 H 3 2 , or HCO2CH3. It is a weak, monobasic 
acid: HCO2CH3 <=t H+ + C0 2 CHr. 



408 COLLEGE CHEMISTRY 

Destructive Distillation of Wood, Charcoal, — Dry wood 
is distilled in iron retorts, and the vapors coming off are led through 
a condenser to separate the liquids from the gases. The cellulose, 
lignin, and resinous material are decomposed, and only charcoal 
remains. The gases, consisting mainly of hydrogen, methane 
CELt, ethane C 2 H 6 , ethylene C2H4, and carbon monoxide CO, are 
employed, on account of their combustibility, as fuel in the dis- 
tillation itself. The fluids form a complex mixture containing 
large quantities of water, methyl alcohol CH 3 OH (wood spirit), 
acetic acid, acetone (CH 3 ) 2 CO, and tar. The liquids can be 
separated. The methyl alcohol (wood spirit) is used in varnish 
making. The acetone has several uses (e.g., p. 394). 

Wood charcoal exhibits the cellular structure of the material 
from which it was made, and is therefore highly porous and has 
an enormous internal surface. When the charcoal is burned, the 
mineral constituents of the wood appear in the ash. This is 
composed of the carbonates of the metallic elements present. 
For certain purposes, charcoals, made in the same fashion as 
the above from bones and from blood, find wide application. 
The former, called bone black, contains much calcium phosphate 
(p. 362). In the old method of making charcoal, which is still 
practised, the wood was piled up, covered with turf, and set on 
fire. All the valuable volatile products were lost, as well as part 
of the charcoal itself. 

' Properties of Charcoal, — Charcoal exhibits certain proper- 
ties which are not shared by other forms of carbon. For example, 
it can take up large quantities of many gases. Boxwood charcoal 
will in this way absorb ninety times its own volume of ammonia, 
fifty-five volumes of hydrogen sulphide or nine volumes of oxygen. 
Freshly made dogwood charcoal (used in making the best gun- 
powder), when pulverized immediately after its preparation, 
often catches fire spontaneously on account of the heat liberated 
by the condensation of oxygen. It is therefore set aside for two 
weeks, to permit the slow absorption of moisture and air. The 
absorbed gases may be removed unchanged by heating the char- 
coal in a vacuum. The phenomenon, described as adsorption, is 
caused by the adhesion of the gases to the very extensive surface 
(due to porosity) which the charcoal possesses. Glass and all other 



THE CARBOHYDRATES AND RELATED SUBSTANCES 409 

solids show the same property, though in a smaller degree (p. 88). 
Solid and liquid bodies are also in many cases taken up by charcoal 
in a similar fashion. Thus, organic dyes, such as indigo, litmus, 
and cochineal, and natural coloring matters (see sugar refining, 
p. 405), which are all more or less colloidal in nature, are removed 
when the liquid is shaken with, or poured through pulverized 
charcoal. The organic materials dissolved in drinking water 
also undergo adsorption in charcoal, but the charcoal soon be- 
comes inactive. Charcoal is likewise used in reducing ores, and 
as a smokeless fuel. 

Coal. — When vegetable matter decomposes, without heating, 
and while covered with sand or clay so that air is excluded, water 
and hydrocarbons are liberated, and the products are peat, bitumi- 
nous coal, or anthracite. 

We are concerned mainly with the products obtained by dis- 
tilling coal, to get coal gas and coke, and with its use as fuel. To 
determine its suitability for various purposes, the coal is analyzed, 
and its heating power is measured. 

In coal analysis, the air-dried material is used. The water is 
determined by heating 1 g. at 105° for 1 hour. Much water lowers 
the fuel value, because heat is wasted in vaporizing it, and in de- 
composing it (cf. p. 386). After re weighing the sample, the coal is 
heated with the Bunsen flame in a covered crucible to drive off 
the volatile matter. After weighing again, air is admitted, and 
strong heat is applied to burn up the fixed carbon (coke). The 
residue is the ash. In the following table the proportions are com- 
pared with seasoned wood on the one hand, and with charcoal and 
coke on the other. 

The calorific power of a coal determines largely its value for 
heating. A sample (about 1 g.) is burned in a bomb calorimeter 
(p. 174). The rise in temperature of a known weight of sur- 
rounding water gives the number of calories. The coal is set on 
fire by a wire heated electrically. Engineers use the number of 
British Thermal Units (1 B.T.U. = heat required to raise 1 lb. 
of water 1° F.) developed by 1 pound of coal. The number of 
B.T.U. = 1.8 X number of calories per gram of coal. Bitumi- 
nous coals give much, and widely varying amounts of volatile 
matter, while anthracite gives very little. The ash is the mineral 



410 



COLLEGE CHEMISTRY 





Water. 


Vola- 
tile 
matter. 


Fixed 
car- 
bon. 


Ash. 


Sul- 
phur. 


Cal. per 


Wood 

Peat 

Bituminous 

Semi-bituminous 

Anthracite 

Charcoal 


20.0 
20.0 
1.3 
4.0 
3.0 
3.2 
2.5 


49.0 

51.6 

36.7 

16.0 

5.6 

4.2 

1.3 


30.0 
25.0 
53.5 
68.5 
80.5 
90.7 
86.3 


1.0 
3.2 
8.5 

11 

10.9 
1.7 

12.4 


6!2 

1.7 
0.5 
0.8 

"i!-3 


3,100 

4,270 
7,800 
7,510 
8,000 
7,580 
7,770 
11,000 


Coke 


Petroleum 



matter of the original plants, with rock material in many speci- 
mens. For coal gas, and even for coke, a coal high in volatile 
matter is chosen. For water gas (p. 386) anthracite or coke is 
employed. 

If the heat of combustion of a coal is known, the amount of 
steam it should furnish can be calculated. It takes 100 cal. to 
raise 1 g. of water from 0° C. to 100° C. and 540 cal. more to 
convert it into steam. If the quantity of steam is too small, the 
furnace, draft, or firing is defective. Too much draft, for ex- 
ample, merely adds additional, useless air to be heated. If the 
flue gas, when analyzed, contains only 3 per cent of carbon dioxide, 
instead of the normal 12 per cent, then for every ton of coal 




Fig. 110. 



burned, 52 tons of unnecessary air were raised to the temperature 
of the furnace. Tests of this kind can control the efficiency of 
every device in a modern factory, and they ought to be in uni- 
versal use. 



Coal Gas. — The gas plant (Fig. 110) includes: (1) The fire- 
brick retorts in which the coal is heated to 1300°, (2) the hydraulic 



THE CARBOHYDRATES AND RELATED SUBSTANCES 411 

main, a wide iron pipe above them in which the tar collects, 

(3) the condenser and wash box for cooling and condensing oils, 

(4) the scrubbers where the ammonia is taken out by water, 

(5) the purifier where hydrogen sulphide is absorbed by hydrated 
ferric oxide and (6) the holder where the gas collects. 

One short ton (2000 lbs.) of the bituminous coal in the above 
table gave: Gas 10,500 cu. ft. with 13 candle power, coke 1325 lbs., 
ammonia 5 lbs. ( = 20 lbs. (NK^SC^ worth $60 per ton), and tar 
12 gallons. The components of the gas were: Illuminants 3.8, 
heating gases 90.2, impurities 6.0. Calorific power of gas 610 
B.T.U. per cu. ft.; sp. gr. (air - 1) 0.43. 

The tar is frequently distilled fractionally and yields: Benzene 
C 6 H 6 , from which aniline and many dyes and drugs are prepared; 
naphthalene Ci H 8 , sold as moth-balls, and the starting point for 
synthetic indigo; anthracene C14H10, from which valuable dyes 
such as alizarin and indanthrene are made; phenol or carbolic 
acid (p. 349), and other useful substances. A rougher separation 
yields tar and pitch, for road-making, preserving timber, and 
waterproofing roofs. 

Coke. — The beehive coke oven is a brick structure shaped like 
a beehive, with an additional opening at the top. The coal which 
fills it burns with a limited supply of air. All the vapors and gas 
burn at the upper opening, and the ammonia and tar and com- 
bustible gas are therefore wasted (cf. p. 340). 

The by-product coke oven is a good deal like a gas plant. The 
chief difference is that the heating is arranged so as to decompose 
as much of the volatile matter as possible, and cause it to leave its 
carbon in the retort. The gas is therefore poor in illuminants, but 
excellent as fuel. The ammonia and tar are diminished in amount, 
but still valuable products. The yield of coke is about 73 per 
cent of the original coal, against 66 per cent from the beehive 
oven. 

Burning coke gives a higher temperature than does coal, be- 
cause no heat is used in vaporizing moisture and volatile matter. 
For the same reason, it burns without flame. Because of these 
and other properties, it is employed in immense quantities in re- 
ducing ores of iron in the blast furnace, as well as for many other 
purposes. 



CHAPTER XXXI 

ORGANIC ACIDS AND SALTS. ALCOHOLS, ESTERS. FOODS 

Thus far, one acid, acetic acid, and two alcohols, methyl and 
ethyl alcohol, have been mentioned. But there are whole series 
of organic acids, corresponding to the series of hydrocarbons. 

Organic Acids and Their Salts, — The general formula of the 
saturated series of monobasic acids is H(C0 2 C n H 2n+ i). Thus: 

Formic acid (w = 0), H(C0 2 H). Palmitic acid (n = 15), H(C0 2 Ci 5 H 3 i). 
Acetic acid (n = 1), H(C0 2 CH 3 ). Stearic acid (n = 17), H(C0 2 C 17 H 3 5). 
Butyric acid [n = 2), H(C0 2 C 2 H 5 ). 

Formic (Lat., an ant) acid is secreted by red ants, and is found 
in stinging nettles. Formic (b.-p. 100.1°), acetic, and butyric 
acids are liquids. Palmitic and stearic acids are solids, and are 
mixed with paraffin in making candles. 

Acids containing relatively less hydrogen are unsaturated. Thus, 
oleic acid (n = 17) is H(C0 2 Ci 7 H 3 3). 

The acids with large molecular weight are insoluble in water. 
All the acids, however, react with sodium hydroxide solution, 
giving soluble salts. Thus, palmitic acid gives sodium palmitate : 

NaOH + H(C0 2 Ci 5 H 3 i) ^±H 2 + Na(C0 2 Ci 5 H 3 i). 

Other salts are sodium formate (p. 385) Na(C0 2 H), sodium acetate 
Na(C0 2 CH 3 ), sodium stearate Na(C0 2 Ci 7 H 35 ), sodium oleate 
Na(C0 2 Ci7H 33 ). Common soap is a mixture of the last two salts 
with sodium palmitate. 

Later, in discussing fats and soap, it will be convenient to abbre- 
viate the formulae. A monobasic acid will be indicated by the 
formula HC0 2 R and a salt by NaC0 2 R or Ca(C0 2 R) 2 , where R 
stands for a hydrocarbon radical or group of atoms, such as 
C n H 2n+ i. In organic chemistry a radical is not always able to 
form an ion. Here the ion is C0 2 R~. 

412 



ALCOHOLS, ESTERS 413 

There are also dibasic acids. Oxalic acid (p. 385) H2C2O4 is the 
simplest of these. It may be made by the oxidation of sugar with 
nitric acid. The white crystals used in the laboratory are the 
hydrate H 2 C 2 4 ,2H 2 0. 

Alcohols. — Methyl alcohol CH 3 OH (p. 408) and ethyl alco- 
hol C2H5OH (p. 406) are the first two members of the series 
C n H 2n +iOH. There are also many alcohols with more than one 
OH group in each molecule. Of these, the one we shall presently 
encounter is glycerine C 3 H 5 (OH)3. The sugars are alcohols with 
several hydroxyl radicals. 

Esters. — When an acid and an alcohol are mixed, an ester 
and water are formed. The action is slow and, being reversible, 
is always incomplete. But, by introduction of a dehydrating 
agent, like concentrated sulphuric acid, the water is removed and 
the change brought to completion. Thus, ethyl alcohol and acetic 
acid, when warmed with sulphuric acid, give ethyl acetate: 

C 2 H 5 OH + HC0 2 CH 3 *± C 2 H 5 C0 2 CH 3 + H 2 0. 

This action has the appearance of a neutralization (p. 256), but is 
different in several ways. Alcohol is not a base, and in aqueous 
solution it does not conduct electricity. Then, true neutralization 
takes place instantly, while the foregoing action, and all like it, 
proceed very slowly. Thus, although acetic acid is a true acid, 
it is not here interacting with a base. 

With the assistance of a dehydrating agent, similar actions take 
place between any alcohol and any acid (organic or inorganic). 
The action of cellulose or of glycerine with nitric acid (p. 350) is 
such an interaction. Again, for example: 

C 3 H 5 (OH) 3 + 3H(C0 2 CH3) ^C 3 H 5 (C0 2 CH 3 ) 3 + 3H 2 0, 
glycerine acetic acid glyceryl acetate 

C 3 H 5 (OH) 3 + 3H(C0 2 C 17 H 35 ) ^C 3 H 5 (C0 2 C 17 H 35 )3 + 3H 2 0. 
glycerine stearic acid glyceryl stearate 

The glyceryl radical C3H5 111 is trivalent, and takes the place of 
three atoms of hydrogen. The products are named as if they 
were salts, but they are not ionized in solution and do not react 
with acids, bases, salts, by instantaneous double decomposition, 



414 COLLEGE CHEMISTRY 

as do true salts. To distinguish them from salts, they are called 
esters — R'(C0 2 R). 

Fats and Animal and Vegetable Oils. — The fats and oils 
found in animal tissue, or pressed from seeds of plants, are com- 
posed mainly of esters. Beef fat is a mixture of about three- 
fourths glyceryl palmitate (palmitin) C 3 H5(C0 2 Ci5H3i)3 and 
glyceryl stearate (stearin) C 3 H5(C0 2 Ci7H35)3, along with one- 
fourth glyceryl oleate (olein) Cs^CC^CnH^. Lard (hog fat) 
contains a much larger proportion of the last (60 per cent) and is 
therefore softer. Butter contains the same esters, along with some 
water and some glyceryl butyrate (butyrin) C3H5(C0 2 C 3 H 7 )3. 
Olive oil contains much olein (75 per cent). Cottonseed oil is 
similar in composition, and is used as a substitute for olive oil, or 
for butter in cooking. 

All these fats and oils contain a certain proportion of the free 
organic acids (see p. 420). These oils must not be confused with 
mineral oils, which are mixtures of hydrocarbons. 

As regards physical properties, these oils are all insoluble in water, 
and the heavier ones also in cold alcohol. They dissolve readily, 
however, in ether, benzene, carbon disulphide, and carbon tetra- 
chloride. Hence, benzene is used in dry cleaning clothing made 
of silk or wool. The two last solvents are used in extracting vege- 
table oils. 

Chemical Properties of Fats and Oils. — All fats and oils, 
when boiled with water, and more rapidly when heated (200°) 
with water in a closed vessel, are decomposed. The ester is 
hydrolyzed, and the actions in the three equations last given 
(p. 413) are reversed. Thus, with stearin: 

C3H 5 (C0 2 C 17 H35)3 + 3H 2 -> C 8 H*(OH) 8 + 3HC0 2 C 17 H 3 5, 
stearin glycerine stearic acid 

and when the mixture is cooled, the acid, being insoluble in water, 
forms a solid cake while the glycerine is in solution in the water. 
If a mixture like beef fat is heated with water in this way, a mix- 
ture of palmitic, stearic, and oleic acids is obtained. The oleic 
acid (liquid) is pressed out, and the residue is mixed with paraffin 
to make candles. The glycerine is separated from the water and 



ALCOHOLS, ESTERS 415 

used in making nitroglycerine (glyceryl nitrate, an ester) and in 
medicine. 

When the fat is heated with aqueous sodium hydroxide, the 
soluble sodium salts of the acids are formed. Since these salts are 
known as soaps, this action of a base on an ester is called saponi- 
fication: 

C 3 H 5 (C0 2 C 17 H35)3 + 3NaOH^C 3 H 5 (OH) 3 + 3Na(C0 2 C 17 H 35 ). 
stearin glycerine sodium stearate 

(a soap) 

When common salt is added to the solution ("salting out"), the 
sodium salts of the three acids (the soap) are coagulated and 
separated as a floating layer, which solidifies when cold. The 
glycerine is contained in the salt solution. 

With potassium hydroxide, the potassium salts are obtained, 
and constitute soft soap. 

The soaps are purified by redissolving and again salting out. 
Dyes and perfumes are often added. Floating varieties are made 
by beating the soap before it solidifies, and so introducing bubbles 
of air. Fine sand or pumice is added to make scouring soaps. 
Mixing with glycerine or sugar gives transparent soap. 

Chemical Properties of Soaps. — Since the soaps are soluble 
salts of sodium, they are largely ionized in solution and interact 
with acids by double decomposition: 

Na(C0 2 C 17 H 35 ) + HC1 -» NaCl + H(C0 2 C 17 H 35 ) I . 

The acids, being insoluble, are precipitated. They also enter into 
double decomposition with other salts. Thus, hard water, con- 
taining compounds of calcium and magnesium in solution, give 
precipitates of the corresponding salts. For example: 

2Na(C0 2 C 17 H 35 ) + CaS0 4 -» Na 2 S0 4 + Ca(C0 2 C 17 H 35 )4 . 

Thus, with hard water, much soap is wasted in precipitating the 
" hardness." 

Colloidal Suspension. — To explain the cleansing power of 
soap, it is necessary to learn more about colloids, for soap in solu- 
tion is essentially colloidal. 



416 COLLEGE CHEMISTRY 

The simplest colloidal suspensions are those of metals like gold 
and platinum. They can be made by forming an electric arc 
between the points of two wires, while the points are immersed in 
water. Liquids of various colors, depending on the degree of 
dispersion (fineness of the particles) of the metal, are thus formed. 
Such a liquid (1) leaves no deposit on filter paper, (2) shows no 
elevation in the boiling-point of the solvent and (3) no depression 
in the freezing-point. (4) The suspended body has little or no 
tendency to diffuse into a layer of the pure solvent. In conse- 
quence, if the colloidal solution is placed in a diffusion-shell, which 
is a test-tube shaped tube of filter-paper or parchment, immersed 
in water, none of the colloid escapes through the pores of the shell. 
Ordinary solutes escape more or less quickly, according to their 
molecular weight. Hence, a diffusion shell can be used to separate 
a mixture of colloidal and non-colloidal material. Thus, salt, if 
present with colloidal starch, or sugar if present with colloidal 
gold, can be removed by changing the water round the shell until 
no more salt (or sugar) is found to come out. This process is 
called dialysis, and was devised by Graham. 

(5) The most striking property of colloids is shown by the 
ultramicroscope. In a perfectly darkened room, a converging 
beam of strong light is sent horizontally 
through the liquid (Fig. Ill) and the place 
where the light is focussed is viewed from 
above, through a microscope. Under such 
circumstances, a true solution remains per- 
fectly dark, but a colloidal suspension shows 
minute points of light, first studied by Tyn- 
dall. Colloidal gold, solutions of soap, starch, 
gelatine and dyes, and many other liquids 
exhibit the phenomenon. The points of light, due to particles 
which, although minute, contain many molecules, show also a 
trembling or vibrating movement, first noticed by a botanist 
Brown (1827) and called the Brownian Movement. The motion 
is due to collisions of the moving molecules of the solvent 
with the suspended particles of the colloid and, when the sus- 
pension is very fine (highly "disperse"), the particles shoot 
about rapidly. 

Other properties of colloidal suspensions are discussed below. 




ALCOHOLS, ESTERS 417 

Theory of Colloidal Suspension and Coagulation. — 

When wires from a battery are immersed in the liquid, the par- 
ticles of a colloid are found to move slowly either with or against 
the positive current. The phenomenon is called electrophoresis. 
Apparently, the colloidal particles are aggregates of molecules of 
an insoluble substance,, collected round an ion. The particles, 
although relatively large in proportion to the charge, move 
almost as rapidly as in ionic migration (p. 231). This affords an 
explanation of the fact that the particles remain suspended, and 
do not settle. They are individually so small that they are kept 
in motion by collisions with the molecules of the solvent. If they 
could unite into large aggregates — like the particles of a precipi- 
tate — they would separate like any ordinary, insoluble substance. 
But, having like electrical charges, they repel one another, and so 
remain separate and in suspension. 

Now those colloids which have distinct electrical charges can 
be coagulated or flocculated, and so precipitated in the liquid, by 
adding a solution of an ionized substance. Thus, colloidal gold 
and other metals are negative, and an equivalent amount of a 
positive ion, usually H + , is present also. When a salt is added, 
the positive ion of the salt attaches itself to the negative colloidal 
metallic particles, neutral bodies result, coagulation can now occur, 
and precipitation follows. Bivalent ions are more effective than 
univalent ones (see arsenic trisulphide) . Conversely, a positive 
colloid, like ferric hydroxide, is coagulated by the negative ion of 
the salt, and more easily the higher the valence of the negative ion. 
Furthermore, one colloid will coagulate another of opposite charge. 
Thus, metaphosphoric acid is a negative colloid when in solution, 
while ortho- and pyrophosphoric acid are not colloidal. Albumin 
is usually a positive colloid. Hence (p. 372), metaphosphoric acid 
and albumin coagulate and precipitate one another, while the other 
two acids have no action on albumin. 

Starch and gelatine are neutral colloids, and are not easily 
coagulated. 

Soap Solution Colloidal, Salting Out. — Soap solution, un- 
der the ultramicroscope, is seen to contain suspended particles. 
A test with litmus also shows that the soap is partly hydrolyzed: 

Na(C0 2 R) + H 2 <± H(C0 2 R) + NaOH. 



418 COLLEGE CHEMISTRY 

Being a salt of a little ionized acid, the negative ion of the soap 
tends to combine with the H + of the water: H + + (C0 2 R)~ ^ 
H(C0 2 R), leaving the ions of sodium hydroxide. Now the acid 
thus set free combines with the undissociated molecules of the salt 
to form an acid salt NaH(C0 2 R)2. This acid salt is insoluble, but 
remains in colloidal suspension as a negative colloid. When a 
strong solution of common salt (or even excess of sodium hydrox- 
ide) is added, the positive ion Na + is adsorbed by the negative 
colloid (the acid salt) and the latter is coagulated. In coming out 
as a precipitate, it seems to adsorb most of the sodium hydroxide, 
so that the precipitate has the composition of soap. 

The Cleansing Power of Soap. Emulsions. — As a cleanser, 
soap solution — or suspension, as we should now call it — has two 
properties. It removes oil and grease (insoluble liquids) by 
forming an emulsion with them. It also removes minute solid 
particles of dirt, by taking the dirt into suspension (next section). 

When an oil, such as kerosene, is violently shaken with water, 
both liquids are broken into minute droplets, and an opaque mix- 
ture results. The droplets of each liquid, however, quickly join 
together and soon the mass clears up and shows the two liquids in 
separate layers. If, however, a colloidal suspension is used, 
instead of pure water, the droplets unite much more slowly, if at 
all, and a more or less permanent, opaque, rather viscous mass 
results. Such a mixture of two mutually insoluble liquids is called 
an emulsion. Thus, a few drops of soap solution will cause the 
kerosene and water to remain much longer in the condition of an 
emulsion. Similarly, vinegar and olive oil, when vigorously 
beaten (French dressing) separate rather quickly into two layers. 
But if the yolk of an egg (colloidal) is added to the vinegar, a stiff, 
almost solid mass can be made (Mayonnaise dressing) which will 
remain permanently emulsified. In removing grease, therefore, 
rubbing with soap solution turns the grease into suspended drop- 
lets (emulsifies it), and so the grease can be washed away. 

This behavior of a colloid can be explained. When the kerosene 
and water are divided into droplets, with a great increase in the 
total surface, and in the surface energy of both, the surface tension 
of water, which is great, favors the reunion of the drops, with 
diminishing surface, and dissipation of the surface energy. Now, 



419 

while ordinary, dilute solutions have a surface tension close to that 
of water, colloidal solutions (such as 0.5 per cent soap solution) 
have a very low surface tension. Hence, the tendency to di- 
minish the surface of droplets of soap solution, by coalescence, is 
slight and ineffective. Furthermore, as predicted by Willard 
Gibbs of Yale University, and proved by experiment, a colloid 
has the peculiarity that it tends to reach a higher concentration 
in the surface layer than it has elsewhere in the liquid. When the 
colloid has adjusted itself to a state of equilibrium, in this regard, 
it resists a decrease in the surface (which would increase its concen- 
tration beyond the equilibrium value), just as much as it resists an 
increase, which would diminish its concentration. The emulsion 
of a colloidal suspension with an immiscible liquid is thus a stable 
condition. 

Experiments confirming this view are easily made. If a solu- 
tion of a dye like methyl violet (colloid) be shaken violently, and 
the froth (large surface in proportion to quantity of liquid) be 
separated, it is found that the liquid produced by the subsidence of 
the froth (an emulsion with air is not permanent), is darker in 
shade, and contains more dye, than an equal amount of the original 
liquid. Soap solution, after being shaken likewise, contains 
relatively more soap in the froth than in the liquid. 

Adsorption of Colloidal Matter. — As we have seen (p. 409), 
when liquids containing colloidal substances, such as dyes and 
natural organic coloring matters, are shaken with pulverized 
charcoal, the colloid is adsorbed by the charcoal — that is, it ad- 
heres to the surface of the particles of the charcoal. This principle 
is used in decolorizing sugar (p. 405) and in " bleaching" oils. 
Xow, soap is also removed by shaking with charcoal or with in- 
fusorial earth, in the same fashion. 

Pulverized charcoal is, relatively, a coarse powder. If soot, 
which is very finely divided carbon, be freed from oil or grease by 
washing with ether, it gives a loose, non-caking powder. If this 
powder be shaken with water, it settles. If it be shaken with 
dilute soap solution, it remains in suspension, and the liquid re- 
sembles ink. The particles are so fine that, instead of carrying 
down the colloidal soap, and forming a precipitate, as charcoal 
does, they attach themselves to the colloidal soap, and remain 



420 COLLEGE CHEMISTRY 

suspended. This is therefore adsorption, with the difference from 
the ordinary phenomenon, that the colloid carries off the adsorb- 
ent, instead of the adsorbent carrying down the colloid. Now 
dirt is composed largely of soot, and equally fine particles of other 
substances. Hence, the soap first emulsifies the oil on the hands 
(or on soiled linen, for example) and then adsorbs the particles of 
dirt which are thus set free. 

Formerly, soap solution was supposed to remove grease (and 
soot?) because of its slight alkaline reaction, due to hydrolysis. 
This explanation must be given up, because: (1) an alkali so 
dilute that it exists in equilibrium with the free fatty acid, can not 
possibly saponify the ester contained in a grease spot. (2) Pure 
alkali of the same concentration (or stronger) has no more emulsi- 
fying power than has water. Such an alkaline solution will 
indeed emulsify an animal or vegetable oil (cod-liver oil, cotton 
oil, castor oil), but it does so by interacting with the free fatty acid 
always present in such oil (p. 414) and forming therefrom a soap. 
Such an alkaline solution does not emulsify kerosene, although 
soap solution does. The emulsifying agency in each case is a soap. 
(3) Very dilute alkali has no more effect upon soot than has water, 
— but soap solution takes clean (greaseless) soot instantly into 
permanent suspension. (4) An aqueous solution of saponin 
C32H54O18, obtainable from several plants, although it contains no 
alkali, lathers, emulsifies, and adsorbs dirt, just as does soap. It 
is a colloid. 

Cyanogen 

Cyanogen CqN 2 . — This compound is prepared by allowing a 
solution of cupric sulphate to trickle into a warm solution of potas- 
sium cyanide. The cupric cyanide, at first precipitated, quickly 
decomposes, giving cuprous cyanide and cyanogen: 

2KNC + CuS0 4 ->Cu(NC)4 + K 2 S0 4 , 
2Cu(NC) 2 -* 2CuNC + C 2 N 2 1 . 

Cyanogen is a very poisonous gas with a characteristic, faint odor. 

Hydrocyanic Acid HNC. — This acid, called also prussic acid, 

is most easily made by the action of an acid upon a cyanide (see 
Potassium cyanide), followed by distillation. It is a colorless liquid 



FOODS 421 

boiling at 26.5°. It has an odor like that of bitter almonds, and is 
highly poisonous. In aqueous solution it is an extremely feeble 
acid. Hydrocyanic acid and the cyanides are unsaturated com- 
pounds, a fact which is illustrated in the two following paragraphs. 

Cyanates and Thiocyanates. — When potassium cyanide is 
fused and stirred with an easily reducible oxide, like lead oxide 
(PbO), the metal (for example, the lead) collects at the bottom of 
the iron crucible in molten form, and potassium cyanate KNCO is 
produced: 

KNC + PbO -> KNCO + Pb. 

Ammonium cyanate NH4NCO, when dissolved in warm water, 
undergoes a profound change. It turns into urea (NH^CO (car- 
bamide), which has the same composition. The former is a salt, 
and is ionized, and enters into double decompositions: the latter 
is not ionized, but is like ammonia, able to combine with acids. 
This is an example of internal rearrangement (p. 148). Two sub- 
stances, of the same composition and molecular weight, but 
different chemical behavior, are called isomers. Urea is one of the 
forms in which nitrogen is eliminated by animals. The prepara- 
tion of what seemed to be a typical product of life from a substance 
(ammonium cyanate) easily made, if necessary, from the four 
elements themselves, was the first case of the artificial production 
of an organic compound, and created a great sensation when it was 
first accomplished by Wohler in 1828. 

When potassium cyanide in aqueous solution is boiled with 
sulphur or with a polysulphide (p. 274), it is converted into potas- 
sium thiocyanate KCNS. This salt, or ammonium thiocyanate 
NH4CNS, is used in testing for ferric ions on account of the deep- 
red color of ferric thiocyanate (c/. p. 182). 

Foods 

Plants and animals contain substances which are similar in 
composition, such as sucrose and lactose (p. 404), starch and 
glycogen (C 6 Hi Os)*, animal fats and vegetable oils (both esters). 
Albumins and other proteins are found in both. They differ, how- 
ever, markedly in the sources of these substances. The plant 
uses simple materials, like carbon dioxide, water, and potassium 



422 



COLLEGE CHEMISTRY 



nitrate. The animal can make no use of these substances — it 
must be fed on complex compounds of animal or vegetable origin. 

Foods, — Since the animal is continuously eliminating carbon 
dioxide, moisture, compounds of nitrogen, salts, and other sub- 
stances, and is also giving off heat, these materials must be re- 
placed, and fuel must be furnished. Like the plant, an animal 
can absorb only dissolved material. But it prepares its own 
solutions in a remarkable laboratory, the digestive tract. The 
production of suitable soluble substances is called digestion. 

The table shows the chief components of animal food, and the 
proportion in which each is present in the chief foods used by man: 



Ash. 



Beef (lean) . . . 

Cod 

Eggs 

Milk * 

Butter 

Cheese (cheddar) 

Oatmeal 

Wheat flour . . . 
Beans (dried) . . 

Almonds 

Maize (green corn) 

Potatoes 

Lettuce 



Water. 


Protein. 


Fat. 


Carbo- 
hydrate. 


73.8 


22.1 


2.9 




82.6 


15.8 


0.4 




73.7 


14.8 


10.5 




87.0 


3.3 


4.0 


5.6 


11.0 


1.0 


85.0 




27.4 


27.7 


36.8 


4.i 


7.3 


16.1 


7.2 


67.5 


11.9 


13.3 


1.5 


72.7 


12.6 


22.5 


1.8 


59.6 


4.8 


21.0 


54.9 


17.3 


75.4 


3.1 


1.1 


19.7 


78.3 


2.2 


0.1 


18.4 


94.7 


1.2 


0.3 


2.9 



1.2 
1.2 
1.0 
0.7 
3.0 
4.0 
1.9 
0.6 
3.5 
2.0 
0.7 
1.0 
0.9 



* The emulsified fat separates slowly as the cream; the protein (casein, colloidally suspended 
in the skim milk) is coagulated by rennet and constitutes cheese; the carbohydrates (lactose, 
a sugar) is then left in the water, along with inorganic salts. 



We observe that the common animal foods, except milk, containing 
lactose (p. 404), carry no carbohydrates (the ox liver contains 
about 2 per cent of glycogen) ; that potatoes and corn, when dried, 
are nearly all starch; that lean meat, dry, is all protein; that some 
seeds (wheat and beans) contain little fat, some (oats) much more 
fat, and some (almonds and nuts) a large amount; and that 
lettuce is mainly water, with useful inorganic salts in solution, and 
cellulose. 

The proteins, several of which have been mentioned (pp. 312, 
340, 359) are white, amorphous substances containing, besides 



foods 423 

carbon, hydrogen, and oxygen, a large proportion of nitrogen 
(16 per cent), some sulphur (1 per cent) and frequently iron and 
phosphorus as well. 

Digestion. — The process of rendering the constituents of food 
soluble is like fermentation (p. 406) — it is performed mainly by 
enzymes. Each class of components is handled by one or more 
enzymes. Thus, starch (in bread and potatoes) is partly digested 
during mastication by ptyalin (an amylase, p. 406) in the saliva, 
and partly by amylopsin in the small intestine. The resulting 
maltose is decomposed into glucose by another enzyme, and passes 
into the blood where it is oxidized, furnishing heat. In diabetes, 
much of the glucose escapes oxidation, and is eliminated. Again, 
the fats are hydrolyzed into the acids and glycerine by lipases 
(fat-splitting enzymes) in the bile, and the acids go into solution 
(probably colloidal). The acids and glycerine recombine to form 
fats in the blood, and are either deposited in the tissues or oxidized. 
Finally, the proteins are changed in a similar way into peptones 
which are soluble in water, and in this form are able to pass through 
the wall of the intestine. 

Fuel Value of Food. — Although food is required to replace 
waste, much of it is needed to furnish energy, by its oxidation, so 
that muscular movements may be maintained, and the tempera- 
ture of the body kept up to its normal value (37° C). Thus, the 
fuel values of foods are important. The average fuel values, ex- 
pressed in large calories (1 Cal. = 1000 cal. as previously denned, 
p. 174), per gram, are: 

Carbohydrates, 4 Cal. Fats, 9 Cal. Proteins, 4 Cal. 

The fuel values per pound (453.6 g.) are 453.6 times greater. 

Healthy life cannot be maintained on one kind of food — a 
mixed diet is necessary. In general, it is held that 100 g. of 
proteins (giving 400 Cal,) per day, and a sufficient amount of other 
foods to give a total fuel value of 2200 cal. is enough for a person 
doing no physical labor. When physical labor is involved, larger 
values, up to 3800 cal. per day, are necessary. The data in the 
table (p. 422) will enable one to calculate the fuel value of 100 g. 
(or of 1 lb.) of each kind of food. 



424 COLLEGE CHEMISTRY 

Exercises. — 1. Make the graphic formulae of ethyl formate, 
ethylene bromide (p. 393), ethyl alcohol. 

2. Make equations for the formation of palmitin (p. 413), the 
saponification of olein (p. 415). 

3. Prepare a summary of the various statements that have been 
made in the text about catalysis (e.g., pp. 29, 59, 156, 160, 279, 
288, 341, 361, 406), and illustrate fully. 

4. Calculate the fuel value of 1 lb. each of (a) oatmeal, (6) 
potatoes, (c) lettuce. 

5. Calculate the weights, both in pounds and in grams, of 100 
Cal. portions of (a) eggs, (b) wheat flour, (c) almonds, (d) lettuce. 

6. At current market prices, what would be the cost per 100 
Cal. portion of beef, cod, butter, and wheat flour, respectively. 



CHAPTER XXXII 

SILICON AND BORON 

In respect to chemical relations there is a close resemblance be- 
tween silicon and carbon. Silicon gives a monoxide, but is quad- 
rivalent in all its other compounds. It is a non-metallic element. 

Occurrence, — Silicon, unlike carbon, is not found in the free 
condition. In combination it is the most plentiful element after 
oxygen, and constitutes more than one-quarter of the crust of the 
earth. The oxide is silica or sand (Si0 2 ), and this oxide and its 
compounds are components of many rocks. In the inorganic 
world silicon is the characteristic element to almost as great an 
extent as is carbon in the organic realm. 

Preparation of Silicon. — When finely powdered magnesium 
and sand are mixed, and one part of the mass is heated, a violent 
action spreads rapidly through the whole: 

2Mg + Si0 2 -> Si + 2MgO. 

At the same time, and especially if excess of the metal is used, 
some magnesium silicide MgzSi is formed also. The mixture is 
treated with a dilute acid which decomposes the magnesium oxide 
and the suicide (see below), and leaves the silicon (amorphous) 
undissolved. When amorphous silicon is dissolved in molten zinc, 
the mass, when solid, contains crystalline silicon. The zinc is 
removed by the action of a dilute acid, the silicon remaining un- 
affected. 

Silicon and ferrosihcon (an alloy of iron and silicon) are now 
made on a large scale, the former by heating sand and carbon, the 
latter by heating a mixture of ferric oxide and sand with carbon 
in the electric furnace (p. 377). 

Properties. — Amorphous silicon is a brown powder. It 
unites with fluorine at the ordinary temperature, with chlorine at 

425 



426 COLLEGE CHEMISTRY 

430°, with bromine at 500°, with oxygen at 400°, with sulphur at 
600°, with nitrogen at about 1000°, and with carbon and boron at 
temperatures attainable only in the electric furnace. It is dis- 
solved by a mixture of hydrofluoric acid and nitric acid, giving 
silicon tetrafluoride. Crystallized silicon forms black needles be- 
longing to the hexagonal system. Silicon and ferrosilicon act 
readily upon a cold solution of sodium hydroxide (cf. p. 56), the 
ortho- or metasilicate of sodium being formed: 

Si + 2NaOH + H 2 -> Na 2 Si0 3 + 2H 2 T. 

This is one of the sources of hydrogen for filling balloons and air- 
ships. 

Silicon Hydride SiH^. — Silicon differs from carbon in giving 
only two well-defined compounds with hydrogen. The chief one 
may be liberated as a gas by the action of hydrochloric acid upon 
magnesium silicide: 

Mg2Si + 4HC1 -> 2MgCl 2 + SiH*. 

The action is like that by which hydrogen sulphide is made. The 
gas is easily inflammable, and burns to form water and silicon 
dioxide. When heated alone, it decomposes into its constituents. 

Silicon Tetrachloride and Tetrafluoride. — The tetrachlo- 
ride SiCLt is formed by direct union of the free elements. It 
may also be prepared by passing chlorine over a strongly heated 
mixture of silicon dioxide and carbon. The products enter a 
condenser in which the tetrachloride assumes the liquid form: 

2C1 2 + Si0 2 + 2C -> SiCU + 2CO. 

Chlorine is unable to displace oxygen from combination with 
silicon, and has, therefore, when alone, no effect upon sand. In 
the above action, therefore, the carbon is used to secure the oxygen 
while the chlorine combines with the silicon. This kind of inter- 
action was formerly used for making many chlorides (e.g., BC1 3 , 
A1C1 3 , CrCl 3 ) from oxides, before simple ways of obtaining the 
elements in the free condition were known. 

Silicon tetrachloride is a colorless liquid (b.-p. 59°), which fumes 
strongly in moist air, and with water precipitates silicic acid: 

SiCU + 4H 2 -» 4HC1 + Si(OH) 4 |. 



silicon 427 

When strong hydrofluoric acid acts upon sand, silicon tetra- 
fluoride SiF 4 is liberated: 

Si0 2 + 4HF -> 2H 2 + SiF 4 . 

Since the water interacts with the tetrafluoride (see below), the 
latter is usually made by heating sand with powdered calcium 
fluoride and excess of sulphuric acid. In this way the hydrogen 
fluoride is generated in contact with the sand, and at the same time 
the sulphuric acid renders the water inactive. Hydrofluoric acid 
acts in a corresponding way upon all silicates (q.v.) } whether these 
are minerals or are artificial silicates like glass (cf. p. 206). 

Silicon tetrafluoride is a colorless gas. It fumes strongly in 
moist air, and acts vigorously upon water. This interaction is 
different from that of the tetrachloride, because the excess of the 
tetrafluoride forms a complex compound with the hydrofluoric 
acid: 

SiF 4 + 4H 2 -» Si(OH) 4 (+ 4HF). (1) 

(4HF) + 2SiF 4 -> 2H 2 SiF 6 . (2) 

3SiF 4 + 4H 2 -» Si(OH) 4 + 2H 2 SiF 6 . 

The silicic acid is precipitated in the water, and may be separated 
by nitration, leaving a solution of hydronuosilicic acid. 

Hydrofluosilicic Acid H 2 SiF Q . — This acid is stable only in 
solution. When the water is removed by evaporation, silicon tet- 
rafluoride is given off, while most of the hydrogen fluoride remains 
to the last. Its salts are decomposed in a corresponding way when 
they are heated. This acid is used in analysis chiefly because its 
potassium salt K 2 SiF 6 is one of the few salts of this metal which 
are relatively insoluble in water. The barium salt is also insoluble, 
but most of the salts of the heavy metals are soluble. 

Silicon Dioxide Si0 2 . — This substance is found in many 
different forms in nature. In large, transparent, six-sided prisms 
with pyramidal ends it is known as quartz or rock crystal. When 
colored by manganese and iron it is called amethyst, when by 
organic matter, smoky quartz. A special arrangement of the 
structure gives cat's eye. Amorphous forms of the same material, 
often colored brown or red with ferric oxide, are agate, jasper, and 
onyx, the last much used in making cameos. Slightly hydrated 



428 COLLEGE CHEMISTRY 

varieties of silica are the opal and flint. Forms produced by or- 
ganisms are sponges and infusorial earth (Tripoli). The latter 
is used in scouring materials and for decolorizing oils (p. 419). 

Silica is found in the hard parts of straw, of some species of 
horsetail (equisetum), and of bamboo. In the form of whetstones 
it is used for grinding. The clear crystals are employed in making 
spectacles and optical instruments and are more transparent to 
ultra-violet light than is glass. Pure sand is used in glass manu- 
facture (q.v.). Recently, small pieces of chemical apparatus have 
been manufactured by fusing quartz (m.-p. 1600°) in the oxy- 
hydrogen flame or the electric furnace. The material does not 
crystallize on cooling, and is amorphous, like glass. Owing to the 
low coefficient of expansion of silica, the vessels can be heated red 
hot and chilled in cold water without risk of fracture. 

Silicates. Water Glass. — Silicon dioxide, although differing 
profoundly from carbon dioxide in its physical nature, nevertheless 
behaves like the latter chemically. Thus, when boiled with 
sodium hydroxide solution it forms sodium metasilicate Na2Si03 
or orthosilicate Na 4 Si04. 

Si0 2 + 2NaOH -> Na 2 Si0 3 + H 2 0. 

The salt is left as a gelatinous solid (" soluble glass") when the 
water is evaporated. The silicates of potassium and sodium may 
also be obtained by boiling sand with the carbonates of these 
metals, or, more rapidly, usually as metasilicates (see below), by 
fusing the mixture: 

Si0 2 + K 2 C0 3 -> K 2 Si0 3 + co 2 t . 

Water glass or soluble glass, being a salt of a feeble acid with an 
active base, gives a solution with an alkaline reaction (p. 271, 383). 
When manufactured for commercial use, it has the composition 
Na 2 Si 2 5 (Na 2 Si0 3 ,Si0 2 ), and gives a less alkaline solution. It is 
used as a filler in cheap soaps, for fireproofing and waterproofing 
timber and textiles, and for preserving eggs. 

Silicic Acid H^SiO^. — When acids are added to a solution of 
sodium silicate, silicic acid is set free. After a little delay it usually 
appears as a gelatinous precipitate. When, however, the silicate 



silicon 429 

is poured into excess of hydrochloric acid, no precipitation occurs. 
The silicic acid remains in colloidal suspension. The acid is ortho- 

Na4Si0 4 + 4HC1 -» 4NaCl + H^iOU, 
Na 2 Si0 3 + 2HC1 + H 2 -» 2NaCl + IWM, 

but the gelatinous precipitate, when dried, loses the elements of 
water. There seem to be no definite stages, indicating the exist- 
ence of various acids, such as we observe with phosphoric acid. 
The final product of drying is the dioxide. Silicic acid is a very 
feeble acid and, therefore, gives no salt with ammonium hydroxide 
(feeble base). 

The suspension of colloidal silicic acid can be freed from the acid 
and sodium chloride (see equation, above) by dialysis (p. 416). 
It is a positive or a negative colloid, according to the mode of 
preparation, and the two kinds are coagulated by addition of salts 
having bivalent negative and positive ions, respectively. 

Mineral Silicates, — While silicic acid is the ortho-acid 
KtSiCXi, and no other silicic acids have been made, the salts are 
most easily classified by imagining them to be derived from various 
acids representing different degrees of hydration of the dioxide 
(c/. p. 369), or, to put it the other way, different degrees of dehy- 
dration of the ortho-acid. The following equations show the 
relation of the ortho-acid to some of the silicic acids whose salts 
are most commonly found amongst minerals: 

H4Si0 4 - H 2 0->H 2 Si0 3 ( = H 2 0,Si0 2 ) Metasilicic acid. 
2H4gi0 4 - H 2 -> H 6 Si 2 7 (- 3H 2 0,2Si0 2 ) ) 
2H4S104 - 3H 2 -> H 2 Si 2 5 (= H 2 0,2Si0 2 ) i 1J1S111C1C acms - 
3H4Si0 4 - 4H 2 -* H&UOs (= 2H 2 0,3Si0 2 ) Trisilicic acid. 

Di- and trisilicates are those derived from acids containing two and 
three units of silicic anhydride, respectively, in the formula. The 
valences of the negative radicals of the acids are shown by the 
number of hydrogen units in the formulae. 

The composition of minerals is often exceedingly complex. This 
is due to the fact that amongst them mixed salts (p. 245) are very 
common, in which the hydrogen of the imaginary acid is displaced 
by two or more metals in such a way that the total quantity of the 



430 COLLEGE CHEMISTRY 

metals is equivalent to the hydrogen. The following list presents 
in tabular form some typical or common minerals arranged accord- 
ing to the foregoing classification: 

ORTHOSILICATES (I^SiO^ METASILICATES (H 2 Si0 3 ) 

Zircon, ZrSi0 4 Wollastonite, CaSi0 3 

Garnet, Ca 3 Al 2 (Si0 4 )3 Beryl, Gl 3 Al 2 (Si0 3 )6 

Mica, KH 2 Al 3 (Si0 4 ) 3 Talc (soapstone), H 2 Mg 3 (Si0 3 ) 4 

Kaolin, H 2 Al 2 (Si0 4 ) 2 ,H 2 Asbestos, Mg 3 Ca(Si0 3 ) 4 

DISILICATE (H 6 Si 2 7 ) TRISILICATE (EL^Os) 

Serpentine, Mg 3 Si 2 7 ,2H 2 Orthoclase (felspar), KAlSi 3 8 

It will be seen that the total valence of the metal units is equal to 
that of the acid radicals. Thus, in beryl there are six equivalents 
of glucinum (beryllium) and six of aluminium, taking the place of 
twelve units of hydrogen in (H 2 Si0 3 ) 6 . 

Mica, which is obtained in large sheets from Farther India, is 
used in making lamp chimneys and as an insulator in electrical 
apparatus. Kaolin, or clay, like mica, is an acid orthosilicate. 
Garnets are pulverized in manganese steel crushers and used in 
making sandpaper. 

Some of these minerals frequently occur mixed together as regu- 
lar components of certain igneous rocks. Thus, granite (p. 2) 
is a more or less coarse mixture of quartz, mica, and felspar. 
Frequently the oblong, flesh-colored or white crystals of the last 
are very conspicuous. Sandstone is composed of sand cemented 
together by clay or by lime, and colored brown or yellow by ferric 
oxide. 

The high melting-point of silica, compared with carbon dioxide, 
and the formation of these complex silicates, indicate that the 
oxide is highly associated (Si0 2 )a-. 

Boron B 

As regards chemical relations, boron, being a uniformly trivalent 
element, is a member of the aluminium family (see Table of periodic 
system). Yet it is a pronounced non-metal, and its oxide and 
hydroxide are acidic: aluminium is a metal, and with its oxide 
and hydroxide basic properties predominate. Boron and its com- 
pounds really resemble carbon and silicon and their compounds in 
all chemical properties, excepting that of valence. 



BORON 431 

Occurrence. — Like silicon, boron is found in oxygen com- 
pounds, namely, in boric acid and its salts. Of the latter, sodium 
tetraborate Na^Oy, or borax, came first from India under the 
name of tincal. It constitutes a large deposit in Borax Lake in 
California. Colemanite, Ca2B 6 0ii,5H 2 0, from California, fur- 
nishes a large part of the commercial supply of compounds of 
boron. 

Preparation. — When boric oxide is heated with powdered 

magnesium (B 2 3 + 3Mg — » 3MgO + 2B), black, amorphous boron 
can be separated with some difficulty from the borides of magnesium 
in the resulting mixture. When excess of powdered aluminium is 
used, hard crystals of boron are formed. 

Properties. — Boron unites with the same elements as does 
silicon (p. 425), but with somewhat greater activity. Like carbon 
(pp. 276, 355), it is also oxidized by hot, concentrated sulphuric or 
nitric acid, the product being boric acid. It interacts with fused 
potassium hydroxide, giving a borate: 

2B + 6KOH -> 2K 3 B0 3 + 3H 2 . 

Boron, when heated with nitrogen, unites to form the nitride BN, 
a white solid. When heated in the electric furnace with carbon, 
it forms a carbide B 6 C. This substance is harder than carborun- 
dum, and stands next to the diamond in this respect (Appendix II). 

The Halides of Boron. — By combined action of carbon and 
chlorine on boric oxide (p. 426), the trichloride of boron BC1 3 may 
be made. It is a liquid (b.-p. 18°) which fumes strongly in moist 
air, and is completely hydrolyzed by water. 

Boron trifluoride BF 3 is made by the interaction of calcium 
fluoride and sulphuric acid with boron trioxide. The mode of 
preparation and the properties of the substance recall silicon 
tetrafluoride (p. 427). It interacts with water, like the latter, 
giving boric acid and hydrofluoboric acid HBF 4 : 

4BF 3 + 3H 2 -> B(OH) 3 + 3HBF 4 . 

Boric Acid and Boron Trioxide. — Boric acid (boracic acid), 
H3BO3 is somewhat volatile with steam, and is found in Tuscany 



432 COLLEGE CHEMISTRY 

in jets of water vapor (soffioni) which issue from the ground. 
Water, retained in basins of brickwork, is placed over the open- 
ings, and from this water, after evaporation, boric acid is obtained 
in crystalline form. As boric acid is a very feeble acid, and withal 
little soluble, it may also be made from sulphuric acid and con- 
centrated borax solution. It crystallizes on cooling the mixture: 

Na 2 B 4 7 + H 2 S0 4 + 5H 2 ±* Na 2 S0 4 + 4H 3 B04 • 

Boric acid crystallizes from water in thin white plates, which are 
unctuous (like graphite and talc) to the touch. Its solubility in 
water is 4 parts in 100 at 19°, and 34 in 100 at 100°. The solution 
scarcely affects litmus. The green tint it confers on the Bunsen 
flame is used as a test for the acid. At 100° the acid slowly loses 
water, leaving metaboric acid HB0 2 , and at 140° tetraboric acid 
is formed: 4HB0 2 — H 2 — > H 2 B 4 7 . Strong heating gives the 
trioxide B 2 3 , a glassy, white solid. When dissolved in water, all 
these dehydrated compounds revert to orthoboric acid H 3 B0 3 . 
The solution of boric acid in water is used as an antiseptic in 
medicine (half-saturated, 2 per cent solution), and sometimes as 
a preservative for milk and other foods. 

Borates, — Borates derived from orthoboric acid are practically 
unknown. The most familiar salt is borax or sodium tetraborate. 
The decahydrate Na 2 B 4 O7,10H 2 O, which crystallizes from water at 
27° in large, transparent prisms, and the pentahydrate which 
crystallizes at 56°, are both marketed. They are made by crystal- 
lization of native borax. In Germany, borax is prepared from 
boracite, found at Stassfurt, by decomposing a solution of the 
mineral with hydrochloric acid: 

MgCl 2 ,2Mg 3 B 8 0i5 + 12HC1 + 18H 2 -» 7MgCl 2 + 16H 3 B0 3 . 

The boric acid is dissolved in boiling water, and sodium carbon- 
ate is added: 4H 3 B0 3 + Na 2 C0 3 -> Na 2 B 4 7 + 6H 2 + C0 2 . In 
California it is made from colemanite by interaction with sodium 
carbonate. 

Since boric acid is a feeble acid, borax is hydrolyzed by water, 
and the solution has a marked alkaline reaction (cf. p. 271). In 
a 0.1N solution (25°), 0.5 per cent is hydrolyzed. 



BORON 433 

When heated with oxides of metals, sodium tetraborate behaves 
like sodium metaphosphate (cf. p. 371), and is used in the form of 
beads in analysis. If its formula be written 2NaB0 2 ,B 2 3 (cf. p. 
369) it will be seen that a considerable excess of the acid anhydride 
is contained in it, and that, therefore, a mixed metaborate may be 
formed by union with some basic oxide. Thus, with a trace of 
cupric oxide, the bead is tinged with blue, from the presence of 
a compound like 2NaB0 2 ,Cu(B0 2 ) 2 . In welding iron, borax is 
scattered on the parts, and combines with the oxide to form a 
fusible mixed borate, which is forced out by the pressure. Borax 
is also mixed with glass in making enamels for cooking utensils. 

Exercises. — 1. Compare and contrast the elements carbon 
and silicon, and their corresponding compounds. 

2. What would be the interaction between aqueous solutions of 
an ammonium salt and of sodium orthosilicate (cf. p. 429)? Why 
is ammonium silicate completely hydrolyzed by water? 



CHAPTER XXXIII 

THE BASE-FORMING ELEMENTS 

In the present chapter a preliminary view of the chemistry of 
the metallic elements is given. 

Physical Properties of the Metals. — Metals show what is 
commonly called a metallic luster, but, as a rule, they do so only 
when in compact form. Magnesium and aluminium exhibit it 
when powdered, but in this condition most metals are black. 

The metals can all be obtained in crystallized form, when a 
fused mass is allowed to cool slowly and the unsolidified portion 
is poured off. In almost all cases the crystals belong to the regu- 
lar system. 

The metals vary in specific gravity from lithium, which is little 
more than half as heavy as water (sp. gr. 0.59), to osmium, whose 
specific gravity is 22.5. Those which have a specific gravity less 
than 5, namely, potassium, sodium, calcium, magnesium, alumin- 
ium, and barium, are called the light metals, and the others the 
heavy metals. 

Most metals are malleable, and can be beaten into thin sheets 
without loss of continuity. Those which are allied to the non- 
metals, however, such as arsenic, antimony, and bismuth, are 
brittle. The order of the elements in respect to this property, 
beginning with the most malleable, is : Au, Ag, Cu, Sn, Pt, Pb, Zn, 
Fe, Ni. 

The tenacity of the metals places them in an order different from 
the above. It is measured by the number of kilograms which a 
piece of the metal 1 sq. mm. in section can sustain without break- 
ing. The values are as follows: Fe 62, Cu 42, Pt 34, Ag 29, Au 27, 
Al 20, Zn 5, Pb 2. 

The hardness is measured by the ease with which the material 
may be scratched by a sharp, hard instrument. Potassium is as 
soft as wax, while chromium is hard enough to cut glass (Ap- 
pendix II). 

434 



THE BASE-FORMING ELEMENTS 



435 



The temperature at which the metal fuses has an important 
bearing on its manufacture. Most of the following melting-points 
are only approximate: 



Mercury 






-39° 


Potassium 






62° 


Sodium 






96° 


Tin . . . 






232° 


Bismuth . 






271° 


Cadmium 






321° 


Lead . . 






327° 



Zinc . . . 
Antimony 
Magnesium 
Aluminium 
Silver . . 
Gold . . . 
Copper . . 



419° 
630° 
651° 
659° 
960° 
1063° 
1083° 



Cast iron . 
Manganese 
Nickel . . 
Chromium 
Iron (pure) 
Platinum . 
Tungsten . 



1150° 
1260° 
1452° 
1520° 
1530° 
1755° 
3267° 



It will be seen that mercury is a liquid, that potassium and 
sodium melt below the boiling-point of water, and that the metals 
down to the foot of the second column can be melted easily with 
the Bunsen flame. 

The methods of manufacture and the treatment of metals are 
much influenced also by their volatility. The following are easily 
distilled: Mercury, b.-p. 357°; potassium and sodium, b. -p. about 
700°; cadmium, b.-p. 770°; zinc, b.-p. 920°. Even the most in- 
volatile metals can be converted into vapor in the electric arc. 

In many cases molten metals dissolve in one another, forming 
alloys. Some alloys are simply solid solutions. Sometimes, as 
in the case of lead and tin, mixtures can be formed in all pro- 
portions. Again, the solubility may be limited, as in the case of 
zinc and lead, where only 1.6 parts of the former dissolve in 100 
parts of the latter. Frequently chemical compounds are formed. 
The colors of alloys are not the average of those of the constitu- 
ents. Thus, the nickel alloy used in coining contains 75 per cent 
of copper and 25 per cent of nickel, yet it shows none of the color 
of the former. 

Alloys in which mercury is one of the components are known 
as amalgams (Gk. fxAXayfm, a soft mass), and are formed with 
especial ease by the lighter metals. Of the common metals, iron 
is tbe least miscible with mercury. 

The good conductivity of metals for electricity distinguishes 
them with some degree of sharpness from the non-metals. They 
show considerable variation amongst themselves, silver conducting 
sixty times^ as well as mercury. The following table gives the 
conductivities of the metals, expressed in terms of the number of 



436 COLLEGE CHEMISTRY 

meters of wire 1 sq. mm. in section which, at 15°, offer a resistance 

of one ohm: 

Silver, cast 62.89 Nickel, cast 7.59 

Copper, commercial . . 57.40 Iron, drawn ..... 7.55 

Gold, cast 46.30 Platinum ...... 5.7-8.4 

Aluminium, commercial 31.52 Steel 5.43 

Zinc, rolled 16.95 Lead 4.56 

Brass 14.17 Mercury 1.049 

To compare these conductivities with those of solutions, it may- 
be said that decinormal hydrochloric acid (p. 240) has a con- 
ductivity on the above scale of 0.035, or a thirtieth of that of 
mercury. 

The world's production (1913) of the metals in metric tons of 
1000 kilos, is approximately^ follows: 



Copper 1,000,000 


Chromium 50,000 


Gold 680 


Zinc 1,000,000 


Nickel 32,000 


Bismuth 500 


Lead 1,000,000 


Silver 7,800 


Cadmium 50 


Tin 120,000 


Tungsten 4,800 


Platinum 9 


Aluminium 79,000 


Mercury 3,000 





General Chemical Relations of the Metallic Elements. — 

Since most of the compounds of the metals are ionogens, their 
solutions, except when the metal is a part of a compound ion, all 
contain the metal in the ionic state, and the resulting substances, 
such as potassium-ion and cupric-ion, have constant properties, 
irrespective of the nature of the negative ion with which they may- 
be mixed. The properties of the ions, simple and compound, are 
much used in making tests in analytical chemistry. On the other 
hand, the chemical properties of the oxides and of the salts in the 
dry state are of importance in connection with metallurgy. 

There are three chemical properties which are characteristic of 
the metallic elements. The first two of them have already been 
discussed somewhat fully. 

1. The metals are able by themselves to form positive radicals 
of salts, and therefore to exist alone as positive ions (pp. 246, 296). 

2. The oxides and hydroxides of the metals are basic (pp. 
94, 296). 

3. Each typical metal has at least one halogen compound which 
is little, if at all, hydrolyzed by water (p. 296). The same thing 
is true of nitrates and other salts involving active acids. 



THE BASE-FORMING ELEMENTS 437 

In reference to the third characteristic, the non-hydrolysis of 
halides of typical metals, a word of explanation is required. Active 
bases (hydroxides of typical metals), such as sodium hydroxide, give, 
with feeble acids, such as H 2 S (p. 271), H3PO4 (p. 370), H 2 C0 3 
(p. 383), H 2 Si0 3 (p. 428), and H3BO3 (p. 432), salts whose solutions 
are alkaline in reaction. This is due to hydrolysis. But active 
bases give, with active acids, such as HC1 and HN0 3 , salts whose 
solutions are neutral in reaction. This is the fact expressed in the 
third characteristic of the metallic elements. The less active bases, 
being hydroxides of less active metallic elements, give, with 
active acids, salts whose solutions are not neutral, but acid in 
reaction. Thus cupric chloride solution is feebly acid. This is 
because there is a tendency for the ions of the water to form the 
slightly dissociated molecules of the base: 

Cu++ + 20H~ + 2H+ -> Cu(OH) 2 + 2H+. 

Finally, a salt derived from a base and an acid, both of which are 
weak, is also hydrolyzed. If the resulting base or acid is insoluble, 
the hydrolysis may go to completion. Aluminium carbonate and 
ammonium silicate (p. 429) are examples of salts which, for this 
reason, are completely hydrolyzed. The resulting mixture may 
have an acid or a basic reaction, if the acid or the base is sufficiently 
soluble and sufficiently active. Thus, ammonium sulphide (NH^S 
solution is alkaline. 

Aside from these points, many features in the behavior of 
metals and their compounds are summed up in the electromotive 
series (p. 260). Before proceeding farther, the reader should re- 
read all the pages referred to above. He should also reexamine 
the various kinds of chemical changes discussed on pp. 166, 197, 
251 et seq. and particularly the varieties of ionic chemical change 
on p. 259. 

Occurrence of the Metals in Nature. — The minerals from 
which metals are extracted are known as ores. They present a 
comparatively small number of different kinds of compounds. 
Most of the metals are found in more than one of these forms, so 
that in the following statement the same metal frequently occurs 
more than once. 

When the metal occurs free in nature, it is said to be native. 




438 COLLEGE CHEMISTRY 

Thus we have gold, silver, metals of the platinum group, copper, 
mercury, bismuth, antimony, and arsenic occurring native (cf. 
p. 60). 

The metals whose oxides are important minerals are iron, man- 
ganese, tin, zinc, copper, and aluminium. The metals are ob- 
tained commercially from the oxides in each of these cases. 

The metals whose sulphides are used as ores are iron, nickel, 
cobalt, antimony, lead, cadmium, zinc, copper, and mercury. 

From the carbonates we obtain iron, lead, zinc, and copper. 
Several other metals, such as manganese, magnesium, barium, 
strontium, and calcium occur in larger or smaller quantities in the 
same form of combination. 

The metals which occur as sulphates are those whose sulphates 
are not freely soluble, namely, lead, barium, strontium, and 
calcium. 

Compounds of metals with the halogens are not so numerous. 
Silver chloride furnishes a limited amount of silver. Sodium and 
potassium chlorides are found in the salt-beds. 

The natural silicates are very numerous, but few are used 
for the preparation of the metals. Many are employed for other 
commercial purposes, kaolin (p. 430) being a conspicuous example. 

Methods of Extraction from the Ores. — The art of extract- 
ing metals from their ores is called metallurgy. Where the metal 
is native, the process is simple, since melting away from the matrix 
(p. 264) is all that is required. Frequently a flux is added. A flux 
usually is a substance which interacts with infusible materials 
to give fusible ones. It combines with the matrix, giving a fusible 
slag (resembling glass). Since the slag is a melted salt, usually 
a silicate, and does not mix at all with the molten metal, separa- 
tion of the products is easily effected. When the ore is a com- 
pound, the metal has to be liberated by our furnishing a material 
capable of combining with the other constituent. The details of 
the process depend on various circumstances. Thus the volatile 
metals, like zinc and mercury, are driven off in the form of vapor, 
and secured by condensation. The involatile metals, like copper 
and iron, run to the bottom of the furnace and are tapped off. 

Where the ore is an oxide it is usually reduced by heating with 
carbon in some form. This holds for the oxides of iron and cop- 



THE BASE-FORMING ELEMENTS 439 

per. Some oxides are not reducible by carbon in an ordinary 
furnace. Such are the oxides of calcium, strontium, barium, mag- 
nesium, aluminium, and the members of the chromium group. 
At the temperature of the electric furnace even these may be re- 
duced, but the carbides are formed under such circumstances, and 
the metals are more easily obtained otherwise. Recently, heat- 
ing the pulverized oxide with finely powdered aluminium has come 
into use, particularly for operations on a small scale. Iron oxide 
is easily reduced by this means, and even the metals manganese 
and chromium may be liberated from their oxides quite readily 
by this action. On account of the great amounts of heat liberated, 
this procedure has received the name aluminothermy (q.v.). 

When the ore is a carbonate, it is first heated strongly to drive out 
the carbon dioxide (cf. p. 381) : FeC0 3 <=± FeO + C0 2 t , and then 
the oxide is treated according to one of the above mentioned 
methods. When the ore is a sulphide, it has to be calcined (p. 
275), in order to remove the sulphur, and the resulting oxide is 
then reduced. 

Chlorides and fluorides of the metals can be decomposed by 
heating with metallic sodium. This method was formerly em- 
ployed in the making of magnesium and aluminium. 

The metals which are not readily secured in any of the above 
ways can be obtained easily by electrolysis of the fused chloride 
or of some other compound. Aluminium is now manufactured 
entirely by the electrolysis of a solution of aluminium oxide in 
molten cryolite. 

Compounds of the Metals: Oxides and Hydroxides. — 

The oxides may be made by direct burning of the metal, by heating 
the nitrates (cf. p. 351), the carbonates (cf. p. 381), or the hydrox- 
ides: Ca(OH) 2 ^CaO + H 2 Ot. They are practically insoluble 
in water, although the oxides of the metals of the alkalies and of 
the metals of the alkaline earths interact with water rapidly to 
give the hydroxides. Oxides are usually stable. Those of gold, 
platinum, mercury, and silver decompose when heated, yet with 
increasing difficulty in this order. The metals, like the non- 
metals, frequently give several different oxides. Those of the 
univalent metals (having the form K 2 0), if we leave cuprous 
oxide and aurous oxide out of account, have the most strongly 



440 COLLEGE CHEMISTRY 

basic qualities. Those of the bivalent metals of the form MgO, 
when this is the only oxide which they furnish, are base-forming. 
Those of the trivalent metals of the form A1 2 3 , known as ses- 
quioxides (Lat. sesqui-, one-half more), are the least basic of the 
basic oxides. The oxides of the forms Sn0 2 , Sb 2 5 , Cr0 3 , and 
Mn 2 7 , in which the metals have valences from 4 to 7, are mainly 
acid-forming oxides, although the same elements usually have 
other lower oxides, which are basic. 

The hydroxides are formed, in the cases of the metals of the 
alkalies and alkaline earths, by direct union of water with the 
oxides. They are produced also by double decomposition when 
a soluble hydroxide acts upon a salt (cf. p. 252). All hydroxides, 
except those of the alkali metals, lose the elements of water when 
heated, and the oxide remains. In some cases the loss takes place 
by stages, just as was the case with orthophosphoric acid (p. 368). 
Thus lead hydroxide Pb(OH) 2 (q.v.) first gives the hydroxide 
Pb20(OH) 2 , then Pb 3 2 (OH) 2 , and then the oxide PbO. The 
hydroxides of mercury and silver, if they are formed at all, are 
evidently unstable, for, when either material is dried, it is found 
to contain nothing but the corresponding oxide. The hydroxides, 
with the exception of those of the metals of the alkalies and 
alkaline earths, are all little soluble in water. 

Compounds of the Metals: Salts. — It may be said, in 
general, that each metal may form a salt by combination with 
each one of the acid radicals. In the succeeding chapters we shall 
describe only those salts which are manufactured commercially, or 
are of special interest for some other reason. The various salts 
will be described under each metal. Here, however, a few re- 
marks may be made about the characteristics of the more com- 
mon groups of salts. 

The chlorides may be made by the direct union of chlorine with 
the metal (cf. p. 160), or by the combined action of carbon and 
chlorine upon the oxide (cf. p. 426). The latter method is used 
in making chromium chloride. The general methods for making 
any salt (p. 146), such as the interaction of a metal with an acid, 
or of the oxide, hydroxide, or another salt with an acid, or the 
double decomposition of two salts, may be used also for making 
chlorides. The chlorides are for the most part soluble in water. 



THE BASE-FORMING ELEMENTS 441 

Silver chloride, mercurous chloride, and cuprous chloride are al- 
most insoluble, however, and lead chloride is not very soluble. 
Most of the chlorides of metals dissolve without decomposition, 
but hydrolysis is conspicuous in the case of the chlorides of the 
trivalent metals, such as aluminium chloride and ferric chloride 
(cf. p. 437). The chlorides of some of the bivalent metals are 
hydrolyzed also, but, as a rule, only when they are heated with 
water. This is the case with the chlorides of magnesium, calcium, 
and zinc. Most of the chlorides are stable when heated, but those 
of the noble metals, particularly gold and platinum, are de- 
composed, and chlorine escapes. The chlorides are usually the 
most volatile of the salts of a given metal, and so are preferred for 
the production of the spectrum of the metal. Some metals, for ex- 
ample iron, form two or more different chlorides. Indium gives 
InCl, InCl 2 , and InCls. 

The sulphides are formed by the direct union of the metal with 
sulphur, or by the action of hydrogen sulphide or of some soluble 
sulphide upon a solution of a salt (cf. p. 273). In one or two cases 
they are made by the reduction of the sulphate with carbon. The 
sulphides, except those of the alkali metals, are but little soluble 
in water. The sulphides of aluminium and chromium are hy- 
drolyzed completely by water, giving the hydroxides, and those 
of the metals of the alkaline earths are partially hydrolyzed (cf. 
p. 273). 

The carbides are usually formed in the electric furnace by inter- 
action of an oxide with carbon (cf. p. 379). Some of them are 
decomposed by contact with water, after the manner of calcium 
carbide, giving a hydroxide and a hydrocarbon. Of this class are 
lithium carbide Li 2 C 2 , barium and strontium carbides BaC 2 and 
SrC 2 , aluminium carbide AI4C3, manganese carbide MnC, and the 
carbides of potassium and glucinum. Others, such as those of 
molybdenum Mo 2 C and chromium Cr 3 C 2 , are not affected by 
water. 

The nitrates may be made by any of the methods used for 
preparing salts. They are all at least fairly soluble in water. 

The sulphates are made by the methods used for making salts, 
and in some cases by the oxidation of sulphides. They are all 
soluble in water, with the exception of those of lead, barium, and 
strontium. Calcium sulphate is meagerly soluble. 



442 COLLEGE CHEMISTRY 






The carbonates are prepared by the methods used for making 
salts. They are all insoluble in water, with the exception of those 
of sodium and potassium. The hydroxides of aluminium and tin 
are so feebly basic that they do not form stable carbonates (cf. 
pp.429, 437). 

The phosphates and silicates are prepared by the methods used 
in making salts. The former are obtained also by special processes 
already described (p. 371). With the exception of the salts of 
sodium and potassium, all the salts of both these classes are 
insoluble. 

For the exact solubilities of a large number of bases and salts 
at 18°, see the Table inside the cover, at the front of this book. 
Solubilities at all temperatures are shown in the diagram, Fig. 58, 
p. 131. 

Exercises. — 1. What do we mean by saying that an oxide is 
strongly or feebly basic, or that it is acidic? 

2. What is meant by the same terms when applied to an 
hydroxide? 

3. Compare the molar solubilities at 18°, (a) of the halides of 
silver, and (b) of the carbonates and (c) oxalates of the metals of 
the alkaline earths, noting the relation between solubility and 
atomic weight. 

4. What is the molar concentration of chloride-ion in saturated 
solutions of silver chloride and lead chloride at 18°, assuming com- 
plete ionization in these very dilute solutions? 



CHAPTER XXXIV 

THE METALLIC ELEMENTS OF THE ALKALIES; 
POTASSIUM AND AMMONIUM 

The metals of this family, with their atomic weights, are: 

Lithium, Li 6.9 Rubidium, Rb 85.5 

Sodium, Xa (Ger. natrium) . 23.0 Caesium, Cs 132.8 

Potassium, K (Ger. kalium) . 39.1 

The Chemical Relations of the Metallic Elements of the 
Alkalies. — The metals which are chemically most active are 
included in this group, and the activity increases with rising atomic 
weight, caesium being the most active positive element of all. A 
freshly cut surface of any of these metals tarnishes by oxidation as 
soon as it is exposed to the air. All of these metals decompose 
water violently (cf. p. 60), liberating hydrogen. The hydroxides 
which are formed by this action are exceedingly active bases, that 
is to say, they give a relatively large concentration of hydroxide-ion 
in solutions of a given molecular concentration (p. 243). In the 
dry form these hydroxides are not decomposed by heating, while 
the hydroxides of all other metals lose water more or less easily. 
In all their compounds the metals of the alkalies are univalent. 

The compounds of ammonium are discussed in connection with 
those of potassium, to which they present the greatest resemblance. 

The solubilities are often decisive factors in connection with the 
preparation and use of salts. The reader will find most of these in 
the table on the inside of the cover, at the front of this book, or in 
the diagram on p. 131, and, as a rule, the values will not be re- 
peated in the descriptive paragraphs. 

Potassium K 

Occurrence. — Silicates containing potassium, such as felspar 
and mica (p. 430), are constant constituents of volcanic rocks. 
These minerals are not used commercially as sources of potassium 

443 



444 COLLEGE CHEMISTRY 

compounds. The salt deposits (see below) contain potassium 
chloride, alone (sylvite) and in combination with other salts, and 
most of the compounds of potassium are manufactured from this 
material. Part of our potassium nitrate, however, is purified 
Bengal saltpeter (p. 347). Potassium sulphate occurs also in the 
salt layers. 

Preparation, — Potassium was first made by Davy (1807) by 
bringing the wires from a battery in contact with a piece of moist 
potassium hydroxide. Globules of the metal appeared at the 
negative wire. Electrolytic processes have now come back into 
use, commercially, molten potassium chloride being the substance 
decomposed. Castner's reduction process involves the heating of 
potassium hydroxide with a spongy mass of carbide of iron (CFe 2 ). 
The potassium passes off as vapor, and is condensed: 
6KOH + 2C -> 2K 2 C0 3 + 3H 2 + 2K. 

Physical and Chemical Properties, — Potassium is a silver- 
white metal (m.-p. 62°). It boils at 720°, giving a green vapor. 

The density of the vapor shows the molecular weight of potas- 
sium to be about, 40, so that the vapor is a monatomic gas. The 
element unites violently with the halogens, sulphur, and oxygen. 
In consequence of the latter fact it is usually kept under petroleum, 
an oil which neither contains oxygen itself, nor dissolves a sufficient 
amount of moisture from the air to permit much oxidation of the 
potassium to take place. A white, crystalline hydride KH is 
formed when hydrogen is passed over potassium heated to 360°. 
When thrown into water it gives potassium hydroxide, and the 
hydrogen is liberated. 

Potassium Chloride KCl, — Sea-water and the waters of salt 
lakes contain a relatively small proportion of potassium com- 
pounds. During the evaporation of such waters, however, the 
potassium compounds tend to accumulate in the mother-liquor 
while sodium chloride is being deposited on the bottom. Hence 
the upper layers of salt deposits are the richest in compounds of 
potassium. Thus, at Stassfurt, near Magdeburg, there is a thick- 
ness of more than a thousand meters of common salt. Above this 
are 25-30 meters of salt layers in which the potassium salts are 
chiefly found. 



POTASSIUM 445 

The chief forms in which potassium chloride is found in the salt 
beds are sylvite KC1 and carnallite KCl,MgCl 2 ,6H 2 0. The latter 
salt is heated with a small amount of water, or with a mother- 
liquor obtained from a previous operation and containing sodium 
and magnesium chlorides. From the clear liquid, when it cools, 
potassium chloride is deposited first and then carnallite. The 
former is taken out and purified, and the latter goes through the 
process again. This potassium chloride is the source from which 
our other potassium compounds are made. It is also our chief 
potassium-bearing fertilizer. It is a white substance crystallizing 
in cubes, melting at about 750°, and slightly volatile at high tem- 
peratures. 

Recently, the giant kelps of the Pacific coast have been used as 
a source of potassium chloride. The dried seaweed contains 9 per 
cent of this salt and about 0.1 per cent of iodine. 

The Other Halides of Potassium. — When iodine is heated 
in a strong solution of potassium hydroxide, potassium iodate and 
potassium iodide are both formed (p. 318) : 

6KOH + 3I 2 -> 5KI + KI0 3 + 3H 2 0. 

The dry residue from evaporation is heated with powdered carbon 
to reduce the iodate, and all the iodide can then be purified by 
recrystallization. The salt forms large, somewhat opaque cubes 
(m.-p. 623°). It is used in medicine and for precipitating silver 
iodide Agl in photography (q.v.). 

The aqueous solution takes up free iodine, forming KI 3 , in 
equilibrium with dissolved iodine: I 3 ~ <=* I~ + I 2 (dslvd). It is 
used in testing for starch, and in reactions in which a solution of free 
iodine is required. 

Potassium bromide KBr may be made in the same way as 
the iodide. It crystallizes in cubes. It is used in medicine and 
for precipitating silver bromide in making photographic plates 
(q.v.). 

The fluoride of potassium K 2 F 2 may be obtained by treating the 
carbonate or hydroxide with hydrofluoric acid. It is a deliques- 
cent, white salt. When treated with an equimolecular quantity 
of hydrofluoric acid it forms potassium-hydrogen fluoride KHF 2 , a 
white salt which is also very soluble. 



446 



COLLEGE CHEMISTRY 



Potassium Hydroxide KOH. — This compound, known also 
as caustic potash, and colloquially as potassium hydrate, was 
formerly made entirely by boiling potassium carbonate with cal- 
cium hydroxide suspended in water (milk of lime) : 

Ca(OH) 2 (solid) <=*Ca(OH) 2 (dslvd) ^20H~+Ca++ ) _^ r m Mdwlw roPn 
K2CO3 <=* 2K++C0 3 = \ ** ^aCOsCdslvd)^ CaCO, 

(solid) . 

The operation is conducted in iron vessels, because porcelain, being 
composed of silicates, interacts with solutions of bases. On 
account of the very limited solubility of the calcium hydroxide 
(0.17 g. in 100 g. Aq), the water takes up fresh portions into solution 
only when the part dissolved has already undergone chemical 
change. The calcium carbonate which is precipitated is, however, 
still more insoluble (0.0013 g. in 100 g. Aq), and hence the action 
goes forward. After the precipitate has settled, the potassium 
hydroxide is obtained by evaporation of the clear liquid, K+ + 
OH - -> KOH. 

Potassium hydroxide is now manufactured by electrolytic 
processes. When a solution of potassium chloride is electrolyzed, 




Fig. 112. 

chlorine is liberated at the anode, and hydrogen and potassium 
hydroxide at the cathode (p. 228). These two sets of products 
must be kept apart, since by their interaction potassium hypo- 
chlorite and potassium chloride would be formed (cf. p. 308). In 
the Castner-Kellner apparatus (Fig. 112), which serves for making 
either potassium or sodium hydroxide, the two end compartments 
are filled with potassium chloride solution (or brine) and contain 
the graphite anodes. The central compartment contains potas- 
sium hydroxide solution and the iron cathode. The positive 



POTASSIUM 447 

current enters by the anodes, and the chlorine is therefore at- 
tracted to and liberated upon the graphite: 2C1~ + 2© — » Cl 2 . 
After rising through the liquid, it is collected for the manufacture 
of liquefied chlorine or of bleaching powder. The ions of potassium 
(or of sodium) are discharged upon a layer of mercury which covers 
the whole floor of the box, and the free metal dissolves in the mer- 
cury, forming an amalgam (p. 435). The layer of mercury extends 
beneath the partitions, and a slight rocking motion given to the 
cell causes the amalgam to flow below the partition into the central 
compartment. Here the potassium leaves the mercury in the 
form of potassium-ion and is attracted by the cathode. Upon 
this, hydrogen from the water is discharged, and the residual 
hydroxide-ion, together with the metal-ion, constitute potassium 
or sodium hydroxide : 

2K+ + 2H+ + 20H" + 20 -> 2K+ + 20H" + H 2 . 

A slow influx of salt solution to the end compartments, and over- 
flow of the alkaline solution in the central cell, are maintained. 
The overflowing liquid contains 20 per cent of the alkali. Since 
there is no undecomposed chloride present in the part of the solu- 
tion which contains the hydroxide, simple evaporation to dryness 
furnishes the solid alkali. Other forms of electrolytic cells, such as 
the Briggs, and the Townsend-Baekeland cells, are also largely 
in use. 

Potassium hydroxide is exceedingly soluble in water, and conse- 
quently, instead of being crystallized from solution, the molten 
residue from evaporation is cast in sticks. The hydroxide is highly 
deliquescent. It also absorbs carbon dioxide from the air, giving 
potassium carbonate. Solutions of the hydroxide have an exceed- 
ingly corrosive action upon the flesh, resolving it into a slimy 
mass by decomposing the proteins. In solution, the base is highly 
ionized, furnishing a high concentration of hydroxide-ion. Com- 
mercially, it is chiefly employed in the making of soft soap. 

Potassium oxide K 2 may be made by heating potassium nitrate 
with potassium in a vessel from which air is excluded : 2KNOs -f- 
10K — * 6K 2 + N 2 . It interacts violently with water, giving the 
hydroxide. When exposed to the air it unites spontaneously with 
oxygen, and a yellow peroxide K 2 4 is formed. The same peroxide 
is formed when potassium burns in air or oxygen. 



448 COLLEGE CHEMISTRY 

Potassium Chlorate KCIG 3 . — The preparation of this salt 
by interaction of potassium chloride with calcium chlorate has 
already been described (p. 313). It is also made by electrolysis of 
potassium chloride solution, the potassium hydroxide and chlorine 
which are liberated being precisely the materials required. All 
that is necessary is to use a warm, concentrated solution and to 
provide for the mixing of the materials generated at the electrodes. 
The salt crystallizes out when the solution cools. 

Potassium chlorate crystallizes in monoclinic plates. It melts 
at about 334°, and at a temperature slightly above this the visible 
liberation of oxygen begins (c/. pp. 27, 29). On account of the 
ease with which its oxygen is liberated, the salt is employed in 
making fireworks and as a component of the heads of Swedish 
matches. It is also used in medicine. 

Potassium perchlorate KC10 4 , formed by the heating of the chlo- 
rate (p. 315), gives white crystals belonging to the rhombic system. 

By adding chlorine-water to potassium carbonate solution, a 
mixture of the chloride and potassium hypochlorite is formed: 

HC1 + HCIO + K 2 C0 3 -> KC1 + KCIO + H 2 + C0 2 . 

The carbonic acid, however, is not completely displaced by the 
HCIO, which is a feeble acid. Hence, the solution is used, under 
the name eau de Javel (often misspelt Javelle), in the household for 
removing stains. 

The mode of preparing potassium bromate KBr0 3 and potassium 
iodate KI0 3 has already been described (p. 318). Potassium 
iodate may be made also very conveniently by melting together 
potassium chlorate and potassium iodide at a low temperature. 
The iodate is much less soluble (see Table) than the chloride, and 
the mixture may be separated by crystallization from water. 

Potassium Nitrate KNO3. — The formation of this salt in 
nature and its mode of extraction and purification have already 
been described (p. 347). This source of supply proved insufficient, 
for the first time, during the Crimean war (1852-55), and a method 
of manufacture from Chile saltpeter (sodium nitrate), which is a 
much cheaper substance, was introduced. Sodium nitrate and 
potassium chloride are heated with very little water, and the sodium 
chloride produced by the action, which is a reversible one, is by 



POTASSIUM 



449 



far the least soluble of the four salts (see Diagram, p. 131). On 
the other hand, in the hot water, the potassium nitrate is by far 
the most soluble. Hence the hot liquid, quickly drained from the 
crystals through canvas, contains the required salt, and most of 
the sodium chloride is in the form of a precipitate. If the solu- 
bility curve of potassium nitrate (p. 131) is examined, it will be 
seen that this salt is but slightly soluble in cold water, and hence 
most of it is deposited when the solution cools. 
The crystals are mixed with little sodium chloride, 
for, as the curve shows, common salt is little less 
soluble at 10° than it is at 100°. 

Potassium nitrate gives long prisms belonging 
to the rhombic system (Fig. 113). It melts at 
about 340°, and when more strongly heated gives 
off oxygen, leaving potassium nitrite (p. 356). 
Although it does not form a hydrate, the crys- 
tals enclose small portions of the mother-liquor, 
and consequently contain both water and im- 
purities. When heated, the crystals fly to pieces n^ _^^^ 
explosively (decrepitate), on account of the vapor- f^\i 3 

ization of this water. Many substances which 
form large crystals and do not melt at a low temperature, behave 
in the same way and for the same reason. In consequence of 
this, the purest salt is made by violent stirring of the solution 
during the operation of crystallization, the result being the forma- 
tion of a crystal-meal. 

Potassium nitrate is used chiefly in the manufacture of gun- 
powder, which contains 75 per cent of the highly purified salt. 
The other components are 10 per cent of sulphur, 14 per cent of 
charcoal, and about 1 per cent of water. The ingredients are 
intimately mixed in the form of paste, and the material when dry 
is broken up and sifted, grains of different sizes being used for 
different purposes. The chemical action which takes place when 
gunpowder is fired in an open space gives chiefly potassium sul- 
phide, carbon dioxide, and nitrogen : 

2KN0 3 + 3C + S -> K 2 S + 3C0 2 + N 2 . 

The explosion occurring in firearms follows a much more complex 
course, and half of the solid product is said to be potassium car- 



450 COLLEGE CHEMISTRY 

bonate (a solid, hence the smoke). One gram yields 264 c.c. of 
gases (0° and 760 mm.), and a much larger volume at the tempera- 
ture of the explosion, and gives 660 calories. The pressure, at the 
temperature of the explosion, if the gases could be confined within 
the volume originally occupied by the gunpowder, would reach 
about forty-four tons per square inch. In recent years, except in 
mining, common gunpowder has been displaced largely by smoke- 
less powder (pp. 358, 359), which, in decomposing, produces no 
solids. 

Potassium nitrate is used also in preserving ham and corned 
beef, on which it confers a red color. 

Potassium Carbonate K 2 C0 8 * — This salt is manufactured 
from potassium chloride, from the Stassfurt deposits. The chlo- 
ride is heated with magnesium carbonate (magnesite), water, and 
carbon dioxide under pressure: 

2KC1 + 3MgC0 3 + C0 2 + 5H 2 -> 2KHMg(C0 3 ) 2 ,4H 2 + MgCl 2 . 

The hydrated mixed salt separates from the liquid containing 
magnesium chloride and is decomposed by heating with water at 
120°. The product is a solution of potassium carbonate, from 
which the precipitated magnesium carbonate is removed by filtra- 
tion and used over again. In some districts potassium carbonate 
is still extracted from wood-ashes, its original source and the origin 
of its name, potash. The sugar beet takes up a considerable 
amount of potash from the soil, and the extract, after removal of 
the sugar, is evaporated and calcined. Wool scourings, when evap- 
orated and calcined, also afford a small supply. 

This salt is usually sold in the form of an anhydrous powder 
(m.-p. over 1000°). When crystallized from water it gives a 
hydrate 2K 2 C0 3 ,3H 2 0. It is deliquescent. Its aqueous solution, 
like that of sodium carbonate (cf. p. 383), has a marked alkaline 
reaction. The commercial name of the substance is pearl ash. 
It is used in making soft soap and hard (i.e., difficultly fusible) 
glass. It is also employed, by interaction with acids, in making 
salts of potassium. 

The use of the bicarbonate KHC0 3 in purifying carbon dioxide 
has already been mentioned (p. 381). Before the nineteenth 
century, this salt was used under the name saleratus (Lat. aerated 



POTASSIUM 451 

salt), a name now sometimes given the baking soda NaHC0 3 
which has displaced it. 

Potassium Cyanide KNC. — This salt is made by heating 
dry potassium ferrocyanide (q.v.) : 

K4Fe(CN) 6 ->4KNC + Fe + 2C + N 2 . 

When the residue is extracted with water, only the potassium 
cyanide dissolves, and it is easily crystallized in pure form from the 
solution. 

Potassium cyanide is extremely soluble in water, and is therefore 
deliquescent. Its poisonous qualities are equal to those of hydro- 
cyanic acid. The acid is so feeble as to be liberated both by the 
moisture and by the carbon dioxide of the air, and hence this salt 
always has an odor of hydrocyanic acid. Potassium cyanide was 
used in electroplating (q.v.), and in extracting gold (q.v.) from 
its ores, but has been displaced by sodium cyanide NaNC, which 
is now less expensive. 

The preparation of potassium cyanate KCNO, a white, easily 
soluble salt, and of potassium thiocyanate KCNS, a white, deli- 
quescent salt, have already been described (p. 421). 

The Sulphate and Bisulphate. — Potassium sulphate K2SO4 
is a constituent of several double salts found in the Stassfurt de- 
posits. It is extracted from schoenite MgS04,K 2 S04,6H 2 and 
kainite MgS0 4 ,MgCl2,K 2 S04,6H 2 0. The former is treated with 
potassium chloride and comparatively little water, whereupon the 
relatively insoluble potassium sulphate crystallizes out, and the 
magnesium chloride remains in the mother-liquor. The crystals 
belong to the rhombic system, contain no water of crystallization, 
and melt at 1066°. This salt is employed in preparing alum (q.v.) 
and is much used as a fertilizer. Since plants take up solutions 
through their cell walls, they can absorb soluble compounds only. 
They are, therefore, dependent, for the potassium compounds 
which they require, upon the weathering out of soluble potassium 
compounds from insoluble silicates containing potassium (p. 430) 
found in the soil. The weathering takes place too slowly to 
furnish a sufficient supply for many crops, particularly that of the 
sugar-beet. Hence potassium sulphate is mixed directly with 
the soil. 



452 COLLEGE CHEMISTRY 

Potassium-hydrogen sulphate (bisulphate) KHS0 4 is made by the 
action of sulphuric acid upon potassium sulphate : K2SO4 + H2SO4 
— » 2KHS0 4 . It crystallizes from water, in which it is very soluble, 
in tabular crystals. Its properties are similar to those of sodium 
bisulphate, which have already been described (p. 288). 

Sulphides of Potassium. — By the treatment of a solution of 
potassium hydroxide with excess of hydrogen sulphide, a solution of 
potassium-hydrogen sulphide is obtained. Evaporation of the solu- 
tion gives a deliquescent, solid hydrate 2KHS,H 2 0. When the 
solution, before evaporation, is treated with an equivalent amount 
of potassium hydroxide, and the water is driven off, potassium 
sulphide K 2 S remains behind (cf. p. 270) : 

KHS + KOH <=> K 2 S + H 2 0. 

Considerable amounts of sulphur can be dissolved in solutions of 
either of these sulphides. By evaporation of the resulting yellow 
liquids, various polysulphides have been obtained. These are 
probably K 2 S 5 , or mixtures of the pentasulphide with K 2 S (cf. p. 
274). Similar substances are produced, as a result of the libera- 
tion and recombination of sulphur, when the solutions are exposed 
to the oxidizing action of the air: 

2KHS + 2 -> 2KOH + 2S. 

Properties of Potassium-ion K + : Analytical Reactions. — 

The positive ionic material of the potassium salts is a colorless 
substance. It unites with all negative ions, and most of the 
resulting compounds are fairly soluble. For its recognition we add 
solutions containing those ions which give with it the least soluble 
salts. Thus, with chloroplatinic acid H 2 PtCl 6 it gives a yellow 
precipitate of potassium chloroplatinate K 2 PtCl6. Since nearly 
one part of this salt dissolves in 100 parts of water, the test is far 
from being a delicate one. Picric acid (p. 349) gives potassium 
picrate KC 6 H 2 (N0 2 ) 3 0, which is much less soluble in water (0.4 
parts in 100 at 15°). Perchloric acid and hydrofluosilicic acid 
likewise give somewhat insoluble salts of potassium. Potassium- 
hydrogen tartrate KHC4H4O6 is precipitated by the addition of 
tartaric acid to a sufficiently concentrated solution of a potassium 



AMMONIUM 453 

salt. The neutral tartrate K0C4H4O6 is much more soluble. The 
latter may be obtained by treating the precipitate with a solution 
of potassium hydroxide. Addition of an acid to this solution 
causes reprecipitation of the bitartrate. 

A much more delicate test for the recognition of a potassium 
compound consists in the examination by means of the spectroscope 
of the light given out by a Bunsen flame, in which a little of the 
salt is held upon a platinum wire. When the amount of potassium 
is considerable, and no other substance which would likewise color 
the flame is present to mask the effect, the violet tint is recognizable 
by the eye. 

Rubidium and Caesium. — In 1860 Bunsen discovered 
several new lines in the spectrum given by materials derived from 
the salts in Durkheim mineral water. Two new elements of the 
alkali group were found to cause their presence, and were named, 
from the colors of the lines which they gave, rubidium (red) and 
caesium (blue). Rubidium is obtainable with relative ease from 
the mother-liquors of the Stassfurt works. 

The metals may be obtained by heating their hydroxides with 
magnesium powder. The hydroxides of these two elements are 
more active as bases than is potassium hydroxide. Their salts are 
very much like those of potassium. 

Ammonium 

The compounds of ammonium claim a place with those of the 
alkali metals because in aqueous solution they give ammonium-ion 
NHi + , a substance which in its behavior closely resembles potas- 
sium-ion. Some of the special properties peculiar to ammonium 
compounds, and particularly the properties of ammonium hydrox- 
ide NH4OH, have been discussed in detail already (pp. 343-345). 

Salts of Ammonium. — Ammonium chloride NH4CI, known 
commercially as salammoniac, like all the other compounds of 
ammonium, is prepared from the ammonia dissolved by the water 
used to wash illuminating gas (p. 411), or that obtained from by- 
product coke ovens (p. 411). It is purified by sublimation, and 
then forms a compact fibrous mass. At 337.8° its vapor exercises 



454 COLLEGE CHEMISTRY 

one atmosphere pressure, and is dissociated into ammonia and 
hydrogen chloride to the extent of 62 per cent (p. 345). 

Ammonium nitrate NH4NO3 is a white crystalline salt which 
may be made by the interaction of ammonium hydroxide and 
nitric acid. When heated gently (m.-p. 166°) it decomposes, giving 
nitrous oxide and water (p. 357). It is used as an ingredient in 
fireworks and explosives. 

When ammonium hydroxide is treated with excess of carbon 
dioxide the solution gives, on evaporation, ammonium bicarbonate 
NH4HCO3. This is a white crystalline salt which is fairly stable 
at the ordinary temperature. It has, however, a faint odor of 
ammonia, and its dissociation becomes very rapid when slight heat 
is applied. When a solution of this salt is treated with ammonium 
hydroxide, the normal carbonate (NEL^CC^ is formed. But this 
salt, when left in an open vessel, loses ammonia very rapidly, and 
leaves the bicarbonate behind. 

Ammonium thiocyanate NH 4 NCS (cf. p. 421) is a white salt 
which finds some application in analysis. 

Ammonium sulphate (NIL^SC^ is a white salt which is used 
chiefly as a fertilizer. By electrolysis of a concentrated solution 
of the bisulphate NH4HSO4, ammonium persulphate (NH4) 2 S208, 
which is less soluble, is formed and crystallizes out (cf. p. 291). 

Solutions of ammonium-hydrogen sulphide NH 4 HS and ammo- 
nium sulphide, (NH^S, made by passing hydrogen sulphide gas into 
ammonium hydroxide, are much used in analysis. The sulphide is 
almost completely hydrolyzed by water into the acid sulphide and 
ammonium hydroxide, its behavior being like that of sodium 
sulphide (p. 271) : 

2NH 3 + H 2 S *± (NILO2S fc* 2NH4+ + S=) <_ tt Q - 

h 2 o±^ oir + H+j-* 11 ^ • 

It is used for the precipitation of sulphides, such as zinc sulphide 
ZnS, which are insoluble in water. Although the S = ions are not 
numerous at any moment, disturbance of the equilibrium by their 
removal, when they pass into combination, causes displacements 
which result in the generation of a continuous supply. The liquid 
smells strongly of ammonia and hydrogen sulphide, on account of 
the dissociation of the parent molecules by reversal of the above 
equilibria. 



AMMONIUM 455 

The solutions, when pure, are colorless. They dissolve free 
sulphur, giving yellow polysulphides similar to those of potassium 
(p. 452). The same yellow substances are also obtained by 
gradual oxidation of ammonium sulphide, when the solution of this 
salt is allowed to stand in a bottle from which the air is imperfectly 
excluded. 

Ammonium Amalgam, — When a salt of ammonium is de- 
composed by electrolysis the NELf*-, upon its discharge, ordinarily 
gives ammonia and hydrogen, and no substance NH4 is obtained. 
If, however, a pool of mercury is used as the negative electrode, the 
NH4 forms an amalgam with it, and there seems to be no doubt that 
this substance is actually present in solution in the mercury. 
While the amalgam is being formed it swells up and gives off the 
decomposition products above mentioned, so that the existence of 
the substance is only temporary. The same material may be 
obtained by putting sodium amalgam into a strong solution of a 
salt of ammonium. The action is a displacement of one ion by 
another (p. 259) : 

Na(dslvd in mercury) + NH4 + — > NH^dslvd in mercury) + Na + . 

This behavior is interesting since it is in harmony with the idea that 
ammonium, if it could be isolated, would have the properties of a 
metal. Substances other than metals are not miscible with 
mercury. 

Ammonium-ion NH^ + : Analytical Reactions, — Ionic am- 
monium is a colorless substance. It unites with negative ions, giv- 
ing salts, which, in the majority of cases, are soluble. Ammonium 
chloroplatinate (NH^PtCle, and to a less extent ammonium- 
hydrogen tartrate NH4HC4H4O6, are insoluble compounds, and 
their precipitation is used as a test. The surest means of recogniz- 
ing ammonium compounds, however, consists in adding a soluble 
base to the substance (cf. p. 345). The ammonium hydroxide, 
which is thus formed, gives off ammonia, and the latter may be 
detected by its odor. 

Exercises. — 1. What kind of metals will, in general, interact 
with solutions of bases (cf. p. 296)? 



456 COLLEGE CHEMISTRY 

2. Why should a mixture of potassium chlorate and antimony 
trisulphide be explosive? 

3. How should you set about making, (a) a borate of potassium, 
(6) potassium pyrophosphate, (c) ammonium nitrite, (d) ammo- 
nium chlorate, (e) ammonium iodide? 



CHAPTER XXXV 

SODIUM AND LITHIUM. IONIC EQUILIBRIUM CONSIDERED 
QUANTITATIVELY 

Sodium chloride forms more than two-thirds of the solid matter 
dissolved in sea-water, and the great salt deposits are largely com- 
posed of it. Sea-plants contain mainly sodium salts of organic 
acids, just as land-plants contain potassium salts. Chile salt- 
peter and albite (a soda feldspar) are important minerals. 

Compounds of sodium are usually cheaper than the correspond- 
ing ones of potassium. Also, since the atomic weight of sodium is 
23, against 39 for potassium, a smaller weight of the sodium com- 
pound will produce the same chemical result. For these two 
reasons, sodium compounds, except in special cases, 
are always used for commercial purposes. 

Preparation. — Sodium was first made by Davy 
(1807) by electrolysis of moist sodium hydroxide. 
It is manufactured by the electrolysis of fused 
sodium hydroxide by a method invented by Castner. 
The negative electrode projects through the bottom 
of the iron vessel containing the fused hydroxide 
(Fig. 114), and here the sodium and hydrogen are 
liberated. This electrode is surrounded by a wire-gauze parti- 
tion to permit circulation of the fused mass, but prevent escape 
of the globules of sodium. This is surmounted by a bell-shaped 
vessel of iron. The positive electrode is an iron cylinder sur- 
rounding the gauze. The sodium and hydrogen liberated at the 
cathode, being lighter than the fused mass, ascend into the iron 
vessel (at A), under the edge of which the hydrogen escapes. 
Oxygen is set free at the anode. The top is closed, to prevent 
the sodium from burning. The melted sodium is ladled into 
molds, like candle molds. 

457 




458 COLLEGE CHEMISTRY 

Properties. — Sodium is a soft, shining metal, melting at 96° 
and boiling at 742°. The green vapor is a monatomic gas. The 
general chemical properties have already been given (p. 443). 
The metal unites with hydrogen to form a hydride NaH, which 
resembles potassium hydride (p. 444). The amalgam with mer- 
cury, when it contains more than a small amount of sodium, is 
solid, and contains one or more compounds of the two elements. 
This amalgam is often used instead of the metal sodium, since 
the dilution and combination with mercury make the interactions 
of the metal more easily controllable. Sodium is used in the 
manufacture of sodium peroxide and of many carbon compounds 
which are used as drugs and dyes. 

Sodium Chloride NaCl. — Common salt is obtained from the 
salt deposits of Stassfurt and Reichenhall (near Salzburg), in 
Cheshire, at Syracuse and Warsaw in New York, at Salina in 
Kansas, in Utah, California, and many other districts. Natural 
brines are obtained from wells in various parts of the world. Since 
the salt can seldom be used directly, on account of impurities which 
it contains, it is purified by recrystallization from water. Natural 
brines, which are sometimes dilute, are often concentrated by 
dripping over extensive ricks composed of twigs. When the re- 
sulting brine is allowed to evaporate slowly by the help of the 
sun's heat, large crystals, sold as "solar salt," are obtained. By 
the use of artificial heat and stirring, smaller crystals of greater 
purity can be secured. In northern Russia, the brine is allowed to 
freeze, and the water thus removed in the form of ice (p. 134). 
Salt intended for table use must be freed from the traces of mag- 
nesium chloride (q.v.) present in the original brine or deposit, for 
this impurity causes it to absorb moisture more vigorously from 
the air. Addition of a little baking soda NaHCC>3 remedies the 
difficulty, by forming the insoluble magnesium carbonate. The 
purest salt for chemical purposes is precipitated from a saturated 
solution of salt by leading into it hydrogen chloride gas. Ex- 
planation of this effect will be given presently (see pp. 466-472). 

Common salt crystallizes in cubes, the faces of which are usually 
hollow. The crystals decrepitate (p. 449) when heated, and melt 
at about 820°. Common salt is the source of all sodium com- 
pounds, with the exception of the nitrate. From it come also 



sodium 459 

most of the chlorine and hydrogen chloride used in commerce. 
It is a necessary article of diet, furnishing, for example, the hydro- 
chloric acid in the gastric juice (p. 147). 

The Hydroxide and Oxides. — Sodium hydroxide NaOH, 

called also, colloquially, caustic soda, is prepared by the action of 
slaked lime upon sodium carbonate, but mainly by the electrolysis 
of a solution of sodium chloride, in both cases precisely as is potas- 
sium hydroxide (p. 446). Sodium hydroxide is a highly deliques- 
cent substance. Its general chemical properties are identical with 
those of potassium hydroxide. It is used in the manufacture of 
soap, in the preparation of paper pulp, and in many other chemical 
industries. 

Sodium peroxide Na^ is made by heating sodium at 300^00° 
in air which has been freed from carbon dioxide. The sodium is 
placed on trays of aluminium, and is passed into the furnace 
against the current of air. In this way, the freshest sodium meets 
the air from which most of the oxygen has been removed, and the 
action is moderated. Conversely, the almost entirely oxidized 
sodium meets the freshest air, and completion of the oxidation is 
thus assured. 

This oxide is the sodium salt of hydrogen peroxide. When 
thrown into water it decomposes in part, in consequence of the 
heat developed, giving sodium hydroxide and oxygen. With care- 
ful cooling, however, much of it can be dissolved. By interaction 
with acids it yields hydrogen peroxide (p. 222). Sodium peroxide 
is now used commercially for oxidizing and bleaching, and, in the 
form of oxone (p. 28), as a source of oxygen. 
The ordinary sodium oxide Na 2 is made 
in the same way as is potassium oxide 
(p. 447). 

The Nitrate and Nitrite. — The occur- 
rence and purification of sodium nitrate 

NaX0 3 have already been described (p. 347). Its crystals are 
of rhombohedral form (Fig. 115). This salt is one of the best of 
fertilizers, since it furnishes to plants the nitrogen which they 
require in a very easily absorbed form. It is used also in the 
manufacture of potassium nitrate, and of nitric acid. 




460 



COLLEGE CHEMISTRY 



Sodium nitrite NaN0 2 is formed by heating sodium nitrate with 
metallic lead and recrystallizing the product (p. 356). 



Manufacture of Sodium Carbonate. — Natural sodium car- 
bonate is found in Egypt and in other parts of the world. At 
Owen's Lake, California, it is secured by solar evaporation of the 
water. The sesquicarbonate Na 2 C03,NaHC03,2H 2 0, being the 
least soluble of the carbonates of sodium, is the .one deposited. 
Locally, small quantities of sodium carbonate are still made by 
the burning of sea-weed. The substance is manufactured from 
sodium chloride in two ways, namely by the Le Blanc process and 
by the Solvay process. In 1900, however, only two factories used 
the former process. 

The Le Blanc process (1791) involves three chemical actions. In 
the first place, sodium chloride is treated with an equivalent 




'&.<&£h'L •."„ 



Fig. 116. 

amount of sulphuric acid in a large cast-iron or earthenware pan. 
The bisulphate thus produced (c/. p. 141), together with the un- 
changed sodium chloride, is raked out on to the hearth of a rever- 
beratory* furnace (Fig. 116), or into a rotating, inclined iron 
cylinder, and heated more strongly until the action is completed: 

NaCl + NaHS0 4 *± Na 2 S0 4 + HClt . 

* So called because the heated gases from the fire are deflected by the roof 
and play upon the materials spread on the bed of the furnace. 



SODIUM 461 

The product of this treatment is called salt cake. The hydrogen 
chloride, which is liberated in both stages, passes through towers 
containing running water in which it is absorbed. The second and 
third actions which follow are conducted in one operation. They 
consist in the reduction of the sodium sulphate by means of 
powdered coal and the interaction of the resulting sulphide of 
sodium with chalk or powdered limestone, leaving finally black ash: 

NaaSCk + 2C -> Na*S + 2C0 2 , 
NasS + CaC0 3 -> Na^COs + CaS. 

Calcium sulphide is not very soluble in water, and is but slowly 
hydrolyzed by it (p. 273), especially when calcium hydroxide is 
present. The sodium carbonate is therefore extracted from the 
black ash by a systematic treatment of the ash with water. The 
ash is placed in a series of vessels at different levels, and a stream 
of water (30-40°) flows from one vessel to another, until, when it 
issues from the last, it is completely saturated with sodium car- 
bonate. When the material in the first of the vessels has been 
exhausted, the water is allowed to enter the second vessel directly, 
and a vessel containing fresh black ash is added at the lower end 
of the series. In this way the most nearly exhausted ash comes in 
contact with pure water, which is in the best position to dissolve 
the remaining sodium carbonate rapidly, while the fresh black ash 
encounters a solution already almost at the point of saturation. 
The commercial survival of the process depends upon the recovery 
of the sulphur from the spent black ash, and of the hydrogen 
chloride. 

The Solvay, or ammonia-soda process (1860), has now displaced 
the Le Blanc process. It differs from the latter by involving 
almost nothing but ionic actions. A solution of salt, containing 
ammonia and warmed to 40°, fills a tower divided by a number of 
perforated partitions. Carbon dioxide, which is forced in below, 
makes its way up through the liquid. The ammonium bicarbonate 
formed by its action undergoes double decomposition with the 
salt, and sodium bicarbonate which is precipitated (sol'ty, 9.6 g. 
in 100 c.c. Aq) settles upon the partitions: 

NaCl + NH4HCO3 *5 NaHC0 3 1 + NH4CI, 
or HCO3" + Na+ *± NaHC0 3 1 . 



462 COLLEGE CHEMISTRY 

The solid sodium bicarbonate, after being freed from the liquid, is 
heated strongly and leaves behind sodium carbonate: 

2NaHC0 3 <- Na 2 C0 3 + H 2 t + C0 2 T. 

The carbon dioxide which is liberated passes through the operation 
once more. The supply of carbon dioxide is generated in lime- 
kilns of special form. The mother-liquor from the sodium bicar- 
bonate contains ammonium chloride. This is decomposed by 
heating with quicklime from the kilns, and the ammonia which is 
thus obtained is available for the treatment of another batch. 

The anhydrous sodium carbonate (soda ash or calcined soda) is 
recrystallized from water, giving the decahydrate Na 2 CO 3 ,10H 2 O, 
soda crystals, or washing soda. The bicarbonate is baking soda. 

Properties of the Carbonate and Bicarbonate, — The com- 
mon form of sodium carbonate consists of large monoclinic crystals 
of the decahydrate Na 2 CO 3 ,10H 2 O. This substance has a fairly 
high aqueous tension, and loses nine of the ten molecules of water 
which it contains when it is exposed in an open vessel (p. 96), 
leaving the monohydrate. When warmed it melts at 35.2°, giving 
a solution of sodium carbonate in water. The deposit from evap- 
oration, above 35.2°, is the monohydrate Na 2 C0 3 ,H 2 0. At higher 
temperatures, or in a dry atmosphere (p. 96), this in turn can be 
completely dehydrated. In aqueous solution, sodium carbonate 
is hydrolyzed (2.3 per cent in 0.1N solution at 25°), and shows a 
marked alkaline reaction (p. 383). The compound is used in large 
amounts for the manufacture of glass and soap, and in the soften- 
ing of water, and is applied in innumerable ways in the scientific 
industries for purposes akin to cleansing. 

All the familiar compounds of sodium, excepting sodium nitrate 
and the peroxide, are made by the treatment of sodium carbonate 
or sodium hydroxide with acids. 

Sodium bicarbonate NaHC0 3 is formed in the Solvay process 
(p. 461). It can be prepared in a state of purity by passing carbon 
dioxide over the decahydrate of sodium carbonate: 

Na 2 CO 3 ,10H 2 O + C0 2 <=> 2NaHC0 3 + 9H 2 0. 

This action is reversible (cf. p. 384), and sodium bicarbonate shows, 
even in the cold, an appreciable tension of carbon dioxide. The 



sodium 463 

aqueous solution of the pure substance is neutral to phenolphthal- 
efn, on account of the small degree of ionization of the ion HC0 3 ~. 
Ordinarily, however, the solution is alkaline, on account of the 
presence of the carbonate, which is hydrolyzed. The-salt is used 
in the manufacture of baking powder and in medicine. 

Baking Powders. — The object of using the powder is to 
generate carbon dioxide in the dough. TjieHDubbles are retained 
because of the presence of the sticky gluten, a protein (p. 3). 
They expand when the dough is heated in baking, and give to the 
bread its porous texture. 

Baking soda, alone, will give off carbon dioxide, but the sodium 
carbonate which it leaves behind has a disagreeable taste and acts 
upon the gluten causing a yellow color and unpleasant smell. It 
also tends to neutralize the acid in the gastric juice and so impedes 
digestion. To prevent this result, sour milk (containing lactic 
acid) and even vinegar are added. Usually, however, a baking 
powder containing an acid substance along with the bicarbonate is 
employed. Potassium bitartrate (cream of tartar) KHC4H4O6 (p. 
452) is most commonly employed, although alum and primary 
sodium or ammonium orthophosphate (p. 370) are also used: 

HKC 4 H40 6 + NaHCOs -> NalTC^HA + H 2 C0 3 -> H 2 + C0 2 . 

The cream of tartar has the advantages that it is somewhat in- 
soluble and does not act noticeably upon the soda before the 
mixing of the dough is complete, and that the sodium-potassium 
tartrate (Rochelle Salts) produced is not harmful. A little starch 
is added to baking powders to keep the particles of the two other 
ingredients apart, and prevent gradual interaction before use. 

For raising bakers' bread, yeast is employed, and time is allowed 
for the propagation of the yeast and its action upon the sugar (p. 
406) in the flour. A little molasses or malt extract is often added, 
to ensure a sufficient supply of sugar. 

The whites of eggs cause cake to rise, largely because they are 
whipped before use, and bubbles of air, which expand when 
heated, are thus introduced. 

Other Salts of Sodium, — Anhydrous sodium sulphate Na 2 S04 
(thenardite) is found in the salt layers. The same salt is contained 



464 COLLEGE CHEMISTRY 

in mineral waters, such as those of Friedrichshall and Karlsbad. 
It is formed in connection with the manufacture of nitric acid from 
sodium nitrate. It is used, as a substitute for sodium carbonate, 
in making inexpensive glass. 

The decahydrate of sodium sulphate Na 2 SO4,10H 2 O (Glauber's 
salt) forms large monoclinic crystals which give up all their water 
of hydration when kept in an open vessel. When heated, the 
crystals melt at 32.4°, giving the sulphate and water. For the 
solubilities of the hydrate and anhydrous substance, see Fig. 59 
(p. 132). 

Sodium thiosulphate Na 2 S203,5H 2 0, formerly called hyposulphite 
of soda, and still called hypo by photographers, is made by boiling 
a solution of sodium sulphite with sulphur (p. 290). A standard 
solution (p. 257) of the thiosulphate is used in determining 
quantities of free iodine: 

2Na 2 S 2 3 + I 2 -* 2NaI + Na 2 S 4 6 . 

Colorless sodium tetrathionate is formed, and the "end point " 
(consumption of all the iodine) can be ascertained by the starch 
test (see p. 480). 

When heated, dry sodium thiosulphate first loses the water of 
hydration, and then decomposes, giving sodium sulphate, which is 
the most stable oxygen-sulphur compound of sodium (c/. p. 290) 
and sodium pentasulphide: 

4Na 2 S 2 3 -> 3Na 2 S0 4 + Na 2 S 5 . 

From the latter, four unit-weights of sulphur can be driven by 
stronger heating. Sodium thiosulphate is used for fixing negatives 
in photography (q.v.), and by bleachers as an antichlor. 

Sodium hyposulphite Na 2 S 2 4 is prepared in solution by the 
action of zinc on sodium bisulphite and excess of sulphurous acid: 

Zn + 2NaHS0 3 + H 2 S0 3 -> Na 2 S 2 4 + ZnS0 3 + 2H 2 0. 

The solution is an active reducing agent, and is employed largely 
by dyers, for example in reducing indigo (insoluble) to indigo white 
(soluble in an alkaline liquid), in preparing the vat of dye. 

Common sodium phosphate is a dodecahydrate of the secondary 
orthophosphate, Na 2 HP0 4 ,12H 2 0. It is made by neutralization 
of phosphoric acid with sodium carbonate. Its properties have 
already been discussed (pp. 370-371). 



sodium 465 

Sodium metaphosphate NaP0 3 is formed in bead tests (p. 371). 

Sodium tetraborate Na2B 4 O7,10H 2 O (borax) forms large, trans- 
parent prisms. When heated it loses water, and leaves the easily 
fusible anhydrous salt in glassy form. Its sources have already 
been discussed under borates (p. 432). It is used as an ingredient 
in glazes for porcelain, in soldering, for bead reactions (p. 433) and 
for preserving food. 

Sodium disilicate Na-jS^Os (cf. p. 428) is used for fireproofing 
wood and other materials, and for preserving eggs. Sand which 
is moistened with it and pressed in molds, forms, after baking, a 
serviceable artificial stone. 

For sodium cyanide, see p. 488. 

Properties of Sodium-ion Na + : Analytical Reactions. — 

Sodium-ion is a colorless ionic material which unites with all 
negative ions. Practically all the salts so formed are soluble in 
water. The only ones which can be precipitated are sodium fluo- 
silicate Na 2 SiF 6 , made by the addition of hydrofluosilicic acid to a 
strong solution of a sodium salt, and sodium-hydrogen pyroanti- 
moniate Na2H 2 Sb20 7 , made by similar addition of the corresponding 
potassium salt. All compounds of sodium confer a yellow color 
on the Bunsen flame, but this test is so delicate that it is shown by 
the traces of sodium contained in almost all substances. 

Lithium. — Lithium occurs in lepidolite (a lithia mica), in 
amblygonite, and in other rare minerals. Traces of compounds of 
the element are found widely diffused in the soil, and are taken up 
by plants, particularly tobacco and beets, in the ashes of which the 
element may be detected spectroscopically. 

The metal is liberated by electrolysis of the fused chloride. The 
specific gravity of the free element (0.53) is lower than that of any 
other metal. Lithium not only floats upon water, but also in the 
petroleum in which it is preserved. 

The metal behaves towards water and oxygen like sodium (p. 50). 
It unites directly and vigorously with hydrogen (LiH), nitrogen 
(LiaN), and oxygen (Li 2 0), forming stable compounds. The rel- 
ative insolubility (see Table) of the hydroxide LiOH, the car- 
bonate Li 2 C0 3 , and the phosphate Li 3 P0 4 ,2H 2 is in sharp 
contrast to the easy solubility of the corresponding compounds of 



466 COLLEGE CHEMISTRY 

the other alkali-metals, and links lithium with magnesium. The 
compounds of lithium give a bright-red color to the Bunsen flame. 
A bright-red and a somewhat less bright orange line are seen in the 
spectrum. The carbonate is used in medicine. 

Ionic Equilibrium, Considered Quantitatively 

In view of the predominance of ionic actions in the chemistry of 
the metals, and of the determinative effect of ionic equilibria on 
many actions, it is essential that we should be prepared in future 
for a more exact consideration of these phenomena than we have 
hitherto attempted. The whole basis for this exact consideration 
has already been supplied, and only more specific application of the 
principles is demanded. The basis referred to, which should now 
be re-read as a preliminary to what follows, is contained in, (1) the 
discussion of chemical equilibrium in general (pp. 177-190), (2) the 
application of the same principles to ionic equilibrium (p. 238), 
and (3) the illustration of this application in the case of cupric 
bromide (pp. 246-251). 

Excess of One Ion. — In the case of cupric bromide, we showed 
that increasing the concentration of the bromide ions displaced the 
equilibrium by favoring the union of the ions to form molecular 
cupric bromide: 2Br~+ Cu++ — > CuBr 2 . This we speak of as a 
repression of the ionization of the cupric bromide. Now, if the sub- 
stance is a slightly ionized one, like a weak acid or a weak base, the 
repression of the ionization through the formation of molecules in 
this way may remove so many of that one of the ions which is not 
present in excess (corresponding to the Cu ++ in the foregoing illus- 
tration), that the mixture will no longer respond to tests for the 
ion so removed. This is an interesting and very common case. 
The behavior of acetic acid, a weak, slightly ionized acid, will serve 
as an illustration. 

In normal solution (60 g. in 1 1.) acetic acid is only 0.004 ionized 
(p. 241), so that, in the equation for the equilibrium, 

(0.996) HC 2 H 3 2 <=> H+ (0.004) + C 2 H 3 2 " (0.004), 

the relative proportions are as shown by the numbers in parenthe- 
sis. If the whole of the acid (60 g.) were- ionized, there would be 
1 g. of hydrogen-ion per liter. Yet, even in the much smaller 



IONIC EQUILIBRIUM, CONSIDERED QUANTITATIVELY 467 

concentration actually present (0.004 g. per liter), the acid taste of 
the H+ and its effect upon indicators can be distinctly recognized. 
If, now, solid sodium acetate is dissolved in the solution, the liquid no 
longer gives an acid reaction with one of the less delicate indicators, 
like methyl orange (p. 258). The explanation is simple. Sodium 
acetate is highly ionized. It gives, therefore, a large concentra- 
tion of acetate-ion to a liquid formerly containing very little. 
This causes a greatly increased union of the H + ions and C 2 H 3 2 ~ 
ions to occur, and the former, being already very few in number, 
disappear almost entirely. Hence the solution becomes, to all 
intents and purposes, neutral. There is no less acetic acid present 
than before, but the concentration of hydrogen-ion is very much 
smaller. 

Formulation and Quantitative Treatment of the Case of 
Excess of One Ion. — If the semi-mathematical mode of formu- 
lating an equilibrium (p. 184), as applied to the case of an ionogen 
(p. 238), be employed here, the foregoing general statements may 
be made more precise and the conclusions clearer. If [H + ] and 
[C2H3O2 - ] represent the molecular concentrations of hydrogen-ion 
and acetate-ion, respectively, and [HC 2 H 3 2 ] that of the acetic 
molecules at equilibrium, then: 

\H+] X [C 2 H 3 Q 2 -] 
[HC 2 H 3 2 ] 

The value of K is constant, whether the strength of the solution of 
acetic acid is great or small, and even when another substance with 
a common ion is present. In the latter case, [C 2 H 3 2 ~] and [H + ] 
stand for the whole concentrations of each of these ionic substances 
from both sources. 

Now, in normal acetic acid [H+] = 0.004, [C 2 H 3 2 ~] = 0.004 (for 
the number of each kind of ions is the same), and [HC 2 H 3 2 ] = 
0.996, practically 1. Substituting in the formula: 
0.004 X 0.004 , g(=a0J6) . 

When, however, sodium acetate is dissolved in the liquid until the 
solution is normal in respect to this substance also, the following 
additional equilibrium has to be considered: 

(0.47) NaC 2 H 3 2 <=* Na+ (0.53) + C 2 H 3 2 " (0.53). 



468 COLLEGE CHEMISTRY 

The concentration of acetate-ion from this source is 0.53, so that, in 
the mixture of acid and salt, the concentration of acetate-ion 
[C 2 H 3 2 ~] will be 0.53 + 0.004 = 0.534, or nearly 134 times larger 
than in the acid alone. Hence, in order that the product [H + ] X 
[C 2 H 3 02~] may recover, as it must, a value much nearer to the old 
one, [H+] must be diminished to something like T ^ ¥ of its former 
magnitude. That is, [H+] will become equal to about 0.00003, 

ttflOMSX 0-534 =g(=a0il6)| 

the rest of the hydrogen-ion uniting with a corresponding amount 
of the acetate-ion to form molecular acetic acid. The effect of 
adding this amount of sodium acetate therefore is, as we have 
seen, to reduce the concentration of the hydrogen-ion below the 
amount which can be detected by use of an indicator like methyl 
orange. 

This effect is of course reciprocal, and the ionization of the 
sodium acetate will be reduced also. But the acetate-ion furnished 
by the acetic acid is relatively so small in amount (0.00003 against 
0.53) that the effect it produces on the ionization of the salt is 
imperceptible. 

It will be noted that the acetate-ion and hydrogen-ion disappear 
in equivalent quantities, for they unite. There is, however, so 
much of the former that the loss it sustains goes unremarked, while 
there is so little of the latter that almost none of it remains. When 
substances of more nearly equal degrees of ionization are used, 
both effects are equally inconspicuous. Thus, sodium chloride and 
hydrogen chloride in normal solutions yield approximately equal 
concentrations of chloride-ion (0.784 and 0.66). Hence, if one 
mole of sodium chloride were to be dissolved in the portion of water 
already containing one mole of hydrogen chloride, the concentra- 
tion of the chloride-ion, at a very rough estimate, would be nearly 
doubled. If this doubling of the concentration of chloride-ion 
almost halved that of the hydrogen-ion (0.784), in order that the 
expression [Cl~] X [H+] -r- [HC1] might remain constant, the 
concentration of the hydrogen-ion would still be about 0.400 and 
therefore 100 times as great as in molar acetic acid. It is thus 
altogether impossible to reduce the concentration of the hydrogen- 
ion given by an active acid like hydrochloric acid below the limit 



IONIC EQUILIBRIUM, CONSIDERED QUANTITATIVELY 469 

at which indicators are affected, for there is no way of introducing 
the enormous concentration of the other ion which the theory 
demands. 

With more crude means of observation than indicators afford, 
effects like this last may sometimes be rendered visible. This was 
the case with cupric bromide solution, to which potassium bromide 
was added (p. 250). The blue of the cupric-ion disappeared from 
view, while much cupric-ion was still present, because the brown 
color of the molecular cupric bromide covered it up completely. 

Special Case of Saturated Solutions, — The commonest as 
well as the most interesting application of the conceptions de- 
veloped above is met with in connection with saturated solutions, 
especially those of relatively insoluble substances. 

The situation in a system consisting of the saturated solution 
and excess of the solute has been discussed already (read p. 127). 
In the case of potassium chlorate, for example, we have the follow- 
ing scheme of equilibria: 

KCIO, (solid) <=± KC10 3 (dslvd) <± K+ + CIO3". 

Solution of the solid is promoted by the solution pressure of the 
molecules, while it is opposed by the osmotic pressure of the dis- 
solved substance, and the solution is saturated when these tenden- 
cies produce equal effects (p. 128). Now it must be noted that the 
tendency directly opposed to the solution pressure is the partial 
osmotic pressure of the dissolved molecules alone. The chief con- 
tents of the solution, the molecules and two kinds of ions of the salt, 
and any foreign material that may be present, are like a mixture of 
gases, and the principle of partial pressure (p. 72) is to be applied. 
The ions and the foreign material do not deposit themselves upon 
the solid, and take, therefore, no part directly in the equilibrium 
which controls solubility. In respect to this the ions are them- 
selves foreign substances. Hence the conclusion may be stated 
that, in solutions saturated at a given temperature by a given solute, 
the concentration of the dissolved molecules of the solute consid- 
ered by themselves will be constant whatever other substances may be 
present. 

The total "solubility" of a substance, as we have used the term 
hitherto, is made up of a molecular and an ionic part. The latter, 



470 



COLLEGE CHEMISTRY 



as we shall presently see, is not constant when a foreign substance 
containing a common ion is already in the liquid. Since the treat- 
ment of the subject requires us now to distinguish between the two 
portions of the solute, a diagram (Fig. 117) will 
assist in emphasizing the distinction. The ma- 
terial at the bottom is the salt. The mole- 
cules and ions are to be thought of as being 
mixed and as being present in numbers repre- 
sented by the factors n and m. Since no foreign 
body is present, the two ions in this case are 
equal in number. 

When we now apply these ideas to the mathe- 
matical expression of the relation: 




Fig. 117. 



[K+] X [CIO3-] 

[KC103] 



= K 



we perceive that, in a saturated solution, [KCIO3], the concentra- 
tion of the molecules, is constant. Transposing, we have 

[K+] X [CIO3-] = X[KC10 3 ] = K'. 

Hence the relation leads to the important conclusion that, in a 
saturated solution, the product of the molar concentrations of the 
ions is constant.* This product is called the ion-product constant 

for the substance. The law of the constancy of the ion-product 
in a saturated solution is one of the most useful of the principles of 
chemistry. It enables us to explain all the varied phenomena of 
precipitation and of the solution of precipitates in a consistent 
manner. These applications of the principle will be explained in 
the next chapter. One curious kind of precipitation will be de- 
scribed here, however, as an illustration of the use of the principle. 



Illustration of the Principle of Ion-Product Constancy. — 

When, to a saturated solution of one of the less soluble salts, a 

* The principle of constant concentration of dissolved molecules, stated 
above, has been shown to express the facts very inaccurately. Now the 
principle of the constancy of the ratio of the ion-product to the concentration 
of the molecules is also inaccurate in the case of highly ionized substances, 
yet in such a way that the two errors neutralize one another. Thus, the 
principle of ion-product constancy here given is in itself fairly exact. 



IONIC EQUILIBRIUM, CONSIDERED QUANTITATIVELY 471 

strong solution of a very soluble salt having one ion in common 
with the first salt is added, precipitation of the first salt frequently 
takes place. This happens, for example, with a saturated solution 
of potassium chlorate, which is not very soluble (molar solubility 
0.52, see Table). The concentrations [K + ] and [C10 3 ~] being small, 
one may easily increase the value for one of the ions, say [C10 3 ~], 
fivefold, by adding a chlorate which is sufficiently soluble. To 
preserve the value of the product [K + ] X [C10 3 ~], the value of 
[K+] will then have to be diminished at once to one-fifth of its 
former value. This can occur only by union of the ionic material 
it represents with an equivalent amount of that for which [C10 3 ~] 
stands. The molecular material so produced will thus tend at first 
to swell the value of [KC10 3 ]. But the value of [KC10 3 ] cannot be 
increased, for the solution is already saturated with molecules, so 
that the new supply of molecules, or others in equal numbers, will 
be precipitated. Hence the ionic part of the dissolved substance 
may be diminished, the equilibria (p. 469) may be partially re- 
versed, and we may actually precipitate a part of the dissolved 
material without introducing any substance, which, in the ordinary 
sense, can interact with it. 

In point of fact, when, to a saturated solution of potassium 
chlorate there is added a saturated solution of potassium chloride 
KC1 (molar solubility, 3.9) or of sodium chlorate NaC10 3 (molar 
solubility, 6.4), a precipitate of potassium chlorate is thrown down. 
These two salts, each containing one of the ions of KCIO3, and being 
much more soluble than the latter (see Table), increase the con- 
centration of one ion and cause the precipitation in the fashion 
just explained. 

The product of the concentrations of the ions, for example [K + ] 
X [C10 3 ~], is called also the solubility product, because these two 
values jointly determine the magnitude of the solubility of the 
substance. The solubility of the molecules is irreducible, but the 
ionic part of the dissolved material may become vanishingly small 
if the value of either [X + ] or [Y~] is very minute. The ionic part 
of any particular substance is made up of the smaller of the two 
concentrations of the ionic substances which it yields, plus an 
equivalent amount, and no more, of the concentration of the other 
ion. The rest of the other ionic substance is part of the solubility 
of some other component. 



472 COLLEGE CHEMISTRY 

Other Illustrations. — The precipitation of sodium chloride 
from a saturated solution, by the introduction of gaseous hydrogen 
chloride (p. 458), is to be explained in the same manner. The 
equilibria : 

NaCl (solid) <=> NaCl (dslvd) <=> Na+ + Cl- 
are reversed by the introduction of additional Cl~ from the very 
soluble, and highly ionized HC1. 

A steady stream of hydrogen chloride is often obtained by drop- 
ping concentrated sulphuric acid into saturated hydrochloric acid: 

H+ + Cl~ <=± HC1 (dslvd) <=> HC1 (gas). 

The effect is due in part to repression of the ionization of the hydro- 
gen chloride and elimination of molecules of the gas from the water 
which is already saturated with molecules of the same kind. 

The formation of potassium hydroxide (p. 446) ceases when a 
certain concentration has been reached. This occurs because the 
concentration of OH~, which rapidly increases, is a factor in the 
solubility product of calcium hydroxide, [Ca ++ ] X [OH - ] 2 . With 
much OH~, little Ca ++ is required to give the constant, numerical 
value of the product. When the concentration [Ca ++ ] from the 
hydroxide has become about as small as that frpm the carbonate, 
the motive for the interaction has been removed. This principle 
is thus as important in industrial operations as it is in analytical 
and other laboratory experimentation. 

Exercises. — 1. The vapor density of sodium peroxide has not 
been determined. Why is the formula Na 2 2 assigned to it? 

2. Construct a scheme of equilibria (p. 271) showing the hy- 
drolysis of calcium sulphide. Why does the presence of calcium 
hydroxide diminish the tendency to hydrolysis (p. 461)? 

3. What will be the effect of adding a concentrated solution of 
silver nitrate to a saturated solution of silver sulphate (see Table 
of solubilities)? 

4. Although a 20 per cent solution of soap can easily be made, 
a 0.5 per cent solution can be salted out (p. 417). How does this 
fact show that salting out is not an operation like the precipitations 
just discussed? 



CHAPTER XXXVI 
THE METALLIC ELEMENTS OF THE ALKALINE EARTHS 

The Chemical Relations of the Elements, — The familiar 
metals of this group, calcium (Ca, at. wt. 40.1), strontium (Sr, at. 
wt. 87.6), and barium (Ba, at. wt. 137.4), constitute a typical 
chemical family, both in the qualitative resemblance to one an- 
other of the elements and of the corresponding compounds, and 
in the quantitative variation in the properties with increasing 
atomic weight. The metals themselves displace hydrogen vigor- 
ously from cold water, giving hydroxides. The solutions of these 
hydroxides, although dilute, on account of a rather small solu- 
bility, are strongly alkaline in reaction. The high degree of ion- 
ization of the hydroxides recalls the hydroxides of the metals of 
the alkalies, and their relative insolubility the hydroxides of the 
"earths" (q.v.). 

In all their compounds, calcium, strontium, and barium are 
bivalent. The hydroxides are formed by union of the oxides with 
water, and are progressively less easy to decompose by heating, 
barium hydroxide being the hardest. The carbonates, when 
heated, yield the oxide of the metal and carbon dioxide, barium 
carbonate being the most difficult to decompose. The nitrates, 
when heated moderately, give the nitrites, but the latter are 
broken up by further heating and yield the oxide of the metal, 
and nitrogen tetroxide. In these and other respects the com- 
pounds of the metals of the alkaline earths resemble those of the 
heavy metals and differ from those of the metals of the alkalies. 
Barium approaches the latter most nearly. 

The table of solubilities (q.v.) shows that the chlorides and 
nitrates of calcium, strontium, and barium are all soluble in 
water, the solubility diminishing in the order given. The sulphates 
and hydroxides cover a wide range from slight solubility to ex- 
treme insolubility. Of the sulphates, 2100, 110, and 2.3 parts, 
respectively, dissolve in one million parts of water. In the case 

473 



474 COLLEGE CHEMISTRY 

of the hydroxides the order of magnitude is reversed, and the cor- 
responding numbers are 200, 630, and 2200. The carbonates are 
almost as insoluble as is barium sulphate. Radium (Ra, at. wt. 
226) belongs to this family (see under Uranium). 

Calcium Ca 

Occurrence. — The fluoride, and the various forms of the car- 
bonate, sulphate, and phosphate, which are found in nature, are 
described below. As silicate, calcium occurs, along with other 
metals, in many minerals and rocks. Compounds of the element 
are found also in plants, and in the bones and shells of animals. 

The Metal. — Calcium is made by electrolysis of the molten 
chloride. A hollow cylinder made of blocks of carbon bolted 
together and open above, forms the anode. A rod of copper 
hanging so that its end dips into the melt forms the cathode. 
The melting of the anhydrous calcium chloride with which the 
cylinder is filled is started by means of a thin rod of carbon laid 
across from the anode to the cathode. When the heat generated 
by the passage of the current through this highly resisting medium 
has melted a sufficient amount of the salt, the rod is removed, and 
the resistance of the fused material suffices to maintain the tem- 
perature. The calcium rises round the cathode and collects on 
the surface of the bath. By slowly elevating the copper cathode, 
the calcium, which adheres to it, may be drawn out of the fused 
mass in the form of a gradually lengthening, irregular rod. The 
rod of calcium is kept constantly in contact with the metal 
which accumulates on the surface, and thus forms one of the 
electrodes. 

Calcium is a silver-white, crystalline metal (m.-p. 800°, sp. gr. 
1.55) which is a little harder than lead, and can be cut, drawn, 
and rolled. It interacts rapidly with water. When heated it 
unites vigorously with hydrogen, oxygen, the halogens, and 
nitrogen. On this account it is used in producing a high degree 
of evacuation. It burns in the air, giving a mixture of the oxide 
and nitride Ca3N 2 . The presence of the latter may be shown by 
the liberation of ammonia when water is added to the residue : 

Ca 3 N 2 + 6H 2 -> 3Ca(OH) 2 + 2NH 3 . 



CALCIUM 475 

A white crystalline hydride CaH 2 is formed by direct union of 
the constituents. It is known in commerce as hydrolyte. It is 
an expensive, but portable source of hydrogen for filling war 
balloons : 

CaH 2 + 2H 2 -> Ca(OH) 2 + 2H 2 . 

Calcium Chloride CaCl 2 * — This salt, for which there is no 
extensive commercial application, is formed as a by-product in 
many industrial operations. Thus, it is a by-product of the 
Solvay soda process (p. 461). By evaporation of any solution, 
the hexahydrate CaCl 2 ,6H 2 is obtained in large, deliquescent, 
six-sided prisms. On account of the great concentration of a 
saturated solution of this compound, the solid and solution do 
not reach a condition of equilibrium with ice (cf. p. 134) until the 
temperature has fallen below —50°. The Solvay process brine 
(p. 462) when mixed with ice, gives, therefore, a very efficient 
freezing mixture. On account of its deliquescent character, the 
solid salt is sprinkled on roads to lay the dust. 

Calcium chloride, partly dehydrated by heating, CaCl 2 ,2H 2 0, 
forms a porous mass which is used in chemical laboratories for 
drying gases and liquids. When complete dehydration is at- 
tempted, the salt interacts with the water, giving some calcium 
oxide. 

Calcium chloride forms compounds, not only with water, but 
also with ammonia (CaCl 2 ,8NH 3 ) and with alcohol. For drying 
these substances, therefore, quicklime is employed. 

Calcium Fluoride CaF 2 . — This compound occurs in nature as 
fluorite or fluor-spar CaF 2 . It crystallizes in cubes, is insoluble in 
water, and when pure is colorless. Natural specimens often 
possess a green tint or show a violet fluorescence. It is formed as a 
precipitate when a soluble fluoride is added to a solution of a salt 
of calcium. 

Fluorite is used in the etching of glass, as the source of the 
hydrogen fluoride (p. 205). It is easily fusible, as its name indi- 
cates (Lat. fluere, to flow), and is employed in metallurgical 
operations as a flux (p. 438), for lowering the melting-point (or 
freezing-point, which is the same thing, cf. p. 134) of the slag (p. 
438), and so facilitating the separation of the latter from the metal,, 



476 COLLEGE CHEMISTRY 

Calcium Carbonate CaCQa, — This compound is found very 
plentifully in nature. Limestone is a compact, indistinctly 
crystalline variety, while marble is a distinctly crystalline form. 
Chalk* is a deposit consisting of the calcareous parts of minute 
organisms. Egg-shells, oyster-shells, coral, and pearls are other 
varieties of organic origin.f Calcite and Iceland spar (Ger. 
spalten, to split) are pure crystallized calcium carbonate. The 
former occurs in flat rhombohedrons, or in pointed, six-sided 
crystals (Fig. 43, p. 83) (scalenohedrons) of "dog-tooth" spar, be- 
longing to the same system. 

When heated, calcium carbonate dissociates, giving carbon dioxide 
and quicklime: 

CaC0 3 ^CaO + C0 2 . 

At ordinary temperatures the decomposition is imperceptible. 
On the contrary, atmospheric carbon dioxide, in spite of its very 
low partial pressure, combines with quicklime, giving "air- 
slaked" lime. As the temperature rises, however, the tension of 
carbon dioxide coming from the carbonate increases, and has a 
fixed value for each temperature. If it is continuously allowed to 
escape, so that the maximum pressure is not reached, the whole 
of the salt eventually decomposes. At 700° the pressure is only 
25 mm., at 900° it reaches an atmosphere, and at 950° two atmos- 
pheres. The phenomenon is precisely similar to the dissociation 
of a hydrate (p. 96) and to the evaporation of a liquid (p. 88). 

Limestone is soluble in water containing carbonic acid, giving 
calcium bicarbonate (p. 384, also see p. 489). By solution of 
limestone, caves are often formed. Conversely, subterranean 
water containing the bicarbonate, when it reaches such a cavern, 
loses carbon dioxide and deposits calcium carbonate as stalactites 
or columns hanging from the ceiling. The drippings form stalag- 
mites on the floors. 

Limestone is used in the manufacture of quicklime (q.v.) and of 
glass. It is employed largely as a flux in metallurgy, when min- 
erals rich in silica are brought into fusible form by the production 
of calcium silicate CaSiOs. Large amounts also find application 
as building-stone. 

* Blackboard " crayon" is usually made of gypsum and not of chalk. 
t The hard coverings of Crustacea and insects are not made of this subs 
stance, bu$ of an organic material called chitin. 



CALCIUM 



477 



Carbon 
""*" dioxide 




Fig. 118. 



Calcium Oxide and Hydroxide. — Pure oxide of calcium CaO 
(quicklime) may be made by ignition of pure marble or calcite. 
For commercial purposes limestone is converted into quicklime 
in kilns (Fig. 118). The flames and heated gases from the fire 
pass between the pieces of limestone, and the carbon dioxide 
liberated is carried off by the draft. 
When the gas is to be used in the 
Solvay process or in the refining of 
sugar, coke (smokeless), instead of 
coal, is employed as the fuel. As 
low a temperature as possible is 
used. A high temperature causes 
impurities in the limestone (e.g., 
clay) to interact with the quick- 
lime, giving fusible silicates, which 
fill the pores and interfere with 
the subsequent slaking with water. - 
Since the change is reversible, if the 
gas lingers in the kiln (at 760 mm. 
pressure), a temperature over 900° 
is required to drive the action forward (p. 476). Hence, a good 
draft, which removes the gas as fast as it is formed, permits the 
use of a lower temperature. 

Pure calcium oxide is a white, porous solid. It is barely fusible 
in the oxyhydrogen flame, but may be melted and boiled in the 
electric arc. It is not reducible by sodium, or by carbon excepting 
at the temperature of the electric furnace. 

When water is poured upon quicklime, it is first absorbed into 
the pores mechanically, and then unites chemically to form calcium 
hydroxide Ca(OH) 2 : 

CaO + H 2 O^Ca(OH) 2 . 

The product is a bulky powder. Much heat is evolved, and part 
of the water is turned into steam. The change is reversible, and 
at a high temperature the hydroxide can be dehydrated. 

Calcium hydroxide is slightly soluble in water: 1 part in 600 
parts of water at 18°, about twice as much water being required at 
100°. The solution, relatively to its concentration, is strongly 
alkaline. On account of its cheapness, this substance is used by 



478 COLLEGE CHEMISTRY 

manufacturers in almost all operations requiring a base, and it 
thus occupies the same position amongst bases that sulphuric acid 
does amongst acids. Caustic lime is employed in the manu- 
facture of alkalies (p. 446), bleaching powder, and mortar (see 
below), the removal of the hair from hides in preparation for 
tanning, the softening of water (see below) and as a whitewash. 

Mortar, — Mortar is made by mixing water with slaked lime 
and a large proportion of sand. The "hardening" process con- 
sists in an interaction of the carbon dioxide of the air with the 
calcium hydroxide: 

C0 2 + Ca(OH) 2 -* CaC0 3 + H 2 0. 

After the superficial parts have been changed, the process goes on 
very slowly, and many years are required before the deeper 
layers have been transformed. The minute crystals of calcite 
which are formed are interlaced with the sand particles, and a 
rigid, yet porous mass is produced. The "hardening" does not 
begin until the excess of water used in making the mortar has 
evaporated, and hence ordinary mortar is unsuitable for use in 
damp places such as cellars. 

Calcium Oxalate CaC 2 4t . — This salt may be observed under 
the microscope in the cells of many plants. It appears in the 
form of needle-shaped or of granular crystals. Since it is the 
least soluble salt of calcium, its formation by precipitation is used 
as a test for calcium ions. 

Theory of Precipitation. — The precipitation of calcium oxa- 
late CaC 2 04, just referred to, is a typical one and may be used to 
illustrate the application of ion-product constancy (p. 470) to 
explaining the phenomenon. The same explanation serves for all 
precipitations of ionogens. 

The first thing to be remembered is that the precipitate which 
we observe, however insoluble its material may be, does not 
include all of the substance, but only the excess beyond what is 
required to saturate the water. The liquid surrounding the pre- 
cipitate is always a saturated solution of the substance precipitated. 
If it were not so, some of the precipitate would dissolve until the 



THEORY OF PRECIPITATION 479 

liquid became saturated. Thus, for example, when we add am- 
monium oxalate solution to calcium chloride solution :j 

<NH4) Sa 5 2cr + cS? 4= I ** CaC2 ° 4 ( dslvd )^ CaC2 ° 4 ( solid ) • 

the liquid is a saturated solution of calcium oxalate, with the excess 
of this salt suspended in it as a precipitate. 

Looking at the matter from this view point, we perceive the ap- 
plication of the rule of ion-product constancy. In this saturated 
solution (p. 470) the product of the ion-concentrations, [Ca^+l X 
[C 2 4 -], is constant. If the original solutions had been so very- 
dilute that, when they were mixed, the product of the concentra- 
tions of these two ions had not reached the value of this constant, 
no precipitation would have occurred. As a matter of fact the ion- 
product considerably exceeded the requisite value, and hence the 
salt was thrown down until the balance remaining gave the value 
in question. The rule for precipitation, then, is as follows: When- 
ever the product of the concentrations of any two ions in a mixture 
exceeds the value of the ion-product in a saturated solution of the 
compound formed by their union, this compound will be precipi- 
tated. Naturally the substances with small solubilities, and there- 
fore small ion-product constants, are the ones most frequently 
formed as precipitates. 

In the case of calcium oxalate, the molar solubility (see Table) 
is O.O443. In so dilute a solution the substance, being a salt 
(p. 242), must be practically all ionized. Each molecule gives 
one ion of each kind. The molar concentrations of these ionic 
substances, Ca 4-4 " and C 2 4 = , in the solution, when the solid is 
also present, must therefore be (practically) 0.0 4 43, each. The 
product [Ca++] X [C 2 4 =] is thus equal to 0.0 4 43 X 0.0 4 43 or 
0.0gl85. If in mixing the solutions, exactly equivalent quan- 
tities were not employed, the values of the two factors will not 
be equal, but the product will in any case possess this value. 

Rule for Solution of Substances, — The rule for solution of 

any ionogen follows at once from the foregoing considerations, and 
may be formulated by changing a few of the words in the rule just 
given: Whenever the product of the concentrations of any two 
ions in a mixture is less than the value of the ion-product in a 



480 COLLEGE CHEMISTRY 

saturated solution of the compound formed by their union, this 
compound, if present in the solid form, will be dissolved. When 
applied to the simplest case, this rule means that a substance will 
dissolve in a liquid not yet saturated with it, but will not dissolve 
in a liquid already saturated with the same material. The value 
of the rule lies in its application to the less simple, but equally 
common cases, such as when an insoluble body is dissolved by 
interaction with another substance (next section). 

Applications of the Rule for Solution to the Solution of 
Insoluble Substances, — So long as a substance remains in pure 
water its solubility is fixed. Thus, with calcium hydroxide, the 
system comes to equilibrium at 18° when 0.17 g. per 100 c.c. of 
water (0.02 moles per liter) have gone into solution: 

Ca(OH) 2 (solid) ±+ Ca(OH) 2 (dslvd) <=> Ca++ + 20H~. 

But if an additional reagent which can combine with either one of 
the ions is added, the concentration of this ion at once becomes less, 
the actual numerical value of the ion-product therefore begins to 
diminish, and further solution must take place to restore its 
value. Thus, if a little of an acid (giving H + ) be added to the 
solution of calcium hydroxide, the union of OH~ and H + to form 
water removes the OH~, and solution of the hydroxide proceeds 
until the acid is used up. There are now more Ca ++ than OH~ 
ions present, but the ion-product reaches the same value as be- 
fore, and then the change ceases. If a further supply of acid is 
added, the removal of OH~ to form H 2 begins again. With 
excess of the acid, the only stable OH~ concentration is that which 
is a factor in the very minute ion-product of water, [OH~] X [H + ], 
which is 0.0 6 1 X 0.0 6 1, or 0.0 13 1. Hence, with excess of acid, the 
calcium hydroxide, which requires in general a much higher con- 
centration of OH - than this to precipitate it or to keep it out of 
solution, finally all dissolves. 

More specifically, if we assume that the calcium hydroxide is 
wholly dissociated in so dilute a solution (which is nearly true), 
each molecule forms one ion of Ca ++ and two ions of OH~. That . 
is, each mole of Ca(OH) 2 gives one mole of Ca ++ and two moles 
of OH~. As the saturated solution contains 0.02 moles of the 
base, the molar concentration (assuming complete dissociation) 



THEORY OF PRECIPITATION 481 

of [Ca++] is 0.02 and of [OH - ] is 0.04. Now, the ion-product is 
the product of the concentrations of all the ions formed, i.e. 
Ca 4-1- , OH - , and OH - . The value of the product is therefore 
[Ca++] X [OH - ] X [OH - ] or [Ca++] X [OH - ] 2 . That is, 0.02 X 
0.04 2 = 0.0 4 32. Note that if the molecule gives two (or three) ions 
of the same kind, the whole concentration of that ion is taken, and 
is also raised to the second (or third) power. 

This particular action is a neutralization of an insoluble base. 
But the other kinds of actions by which insoluble ionogens pass 
into solution all resemble it closely, and differ only in details. The 
general outlines of the explanation are the same in every case. 
We proceed now to apply it to the common phenomenon of the 
solution of an insoluble salt by an acid. 

Interaction of Insoluble Salts with Acids, Resulting in 
Solution of the Salt. — Calcium oxalate passes into solution 
when in contact with acids, especially active acids. Thus, with 
hydrochloric acid, it gives calcium chloride and oxalic acid, both 
of which are soluble: 

CaC 2 4 T + 2HC1 fc; CaCl 2 + H 2 C 2 4 . (1) 

The action of acids upon insoluble salts is so frequently mentioned 
in chemistry and is so important a factor in analytical operations 
that it demands separate discussion. This example is a typical 
one and may be used as an illustration. 

According to the rules already explained (p. 479), calcium 
oxalate (or any other salt) is precipitated when the numerical value 
of the product of the concentrations of the two requisite ions 
[Ca" 1-1 "] X [C 2 4 = ] exceeds the value of the ion-product for a 
saturated solution of calcium oxalate in pure water, that is, ex- 
ceeds 0.0 8 185 (p. 479). When, on the contrary, the product of 
the concentrations of the two ions falls below the limiting value, 
a condition which may arise from the removal in some way either 
of the Ca 4-1- or of the C 2 4 = ions, the undissociated molecules will 
ionize, and the solid will dissolve to replace them until the ionic 
concentrations necessary for equilibrium with the molecules have 
been restored or until the whole of the solid present is consumed. 
Here the oxalate-ion from the calcium oxalate combines with the 



482 COLLEGE CHEMISTRY 

hydrogen-ion of the acid (usually an active one) which has been 
added, and forms molecular oxalic acid: 

C 2 4 = + 2H+ £+ H0C2O4. (2) 

Hence, dissociation of the dissolved molecules of calcium oxalate 
proceeds, being no longer balanced by encounters and unions of 
the now depleted ions, and this dissociation in turn leads to solu- 
tion of other molecules from the precipitate. 

It will be seen that the removal of the ions in this fashion can 
result in considerable solution of the salt only when the acid pro- 
duced is a feebly ionized one. Here, to be specific, the concentra- 
tion of the C 2 4 =:: in the oxalic acid equilibrium, (2) above, must be 
less than that of the same ion in a saturated calcium oxalate 
solution. Now oxalic acid does not belong to the least active 
class of acids, and its pure solution contains a considerable con- 
centration of C 2 04 = . There is, however, a decisive factor in the 
situation which we have not yet taken into account. The hydro- 
chloric acid which we used for dissolving the precipitate is a very 
highly ionized acid and gives an enormously greater concentration 
of hydrogen-ion than does oxalic acid. Hence the hydrogen-ion 
is in excess in equation (2), and the condition for equilibrium, 

r-rr r^ <-\i^ = Kf wm ^ e satisfied by a correspondingly small 

[M2C2U4J 

concentration of C2O4—. In this particular case, therefore, the 
[C 2 04 = ] of the oxalic acid is less than that given by the calcium 
oxalate. The whole change, therefore, depends for its accomplish- 
ment, not only on the mere presence of hydrogen-ion, but on the 
repression of the ionization of the oxalic acid by the great excess of 
hydrogen-ion furnished by the active acid that has been used. As 
a matter of fact, we find that a weak acid like acetic acid has 
scarcely any effect upon a precipitate of calcium oxalate. An 
acid stronger than oxalic acid must be employed. The whole 
scheme of the equilibria is as follows: 

CaC 2 4 (solid)<=»CaC 2 4 (dslvd)^Ca+++C 2 4 =) ^ r n , AaUrA , 
2HC1 *=;2C1-+2H+ j^-^^U (aslvd). 

When excess of an acid sufficiently active to furnish a large con- 
centration of hydrogen-ion is employed, the last equilibrium is 
then driven forward and the others follow. With addition of a 



THEORY OF PRECIPITATION 483 

weak acid, only a slight displacement occurs, and the system comes 
to rest again when the molecular oxalic acid has reached a sufficient 
concentration. 

A generalization may be based on these considerations : an insoluble 
salt of a given acid will in general interact and dissolve when treated 
with a solution containing another acid, provided that the latter acid 
is a much more highly ionized (more active) one than the former 
(see below). 

But even active acids frequently fail to bring salts of weak acids 
into solution, especially when the weak acid is itself present also. 
Here the cause lies in the fact that such salts are even less soluble 
than those of the calcium oxalate type, and give so low a con- 
centration of the negative ion that the utmost repression of the 
ionization of the corresponding acid does not give a lower value 
for the concentration of this ion than does the salt itself. Thus, 
we have seen (p. 272) that even hydrochloric acid (dilute) will not 
dissolve a number of sulphides. For example, in the case of 
cupric sulphide in a solution saturated with hydrogen sulphide, 
the S— factor in the solubility product [Cu ++ ] X [S— ] remains 
smaller than that in the scheme defining the hydrogen sulphide 
equilibrium [H + ] 2 X [S = ] even when the [S = ] factor in the latter 
is diminished in consequence of great addition of hydrogen-ion. 
In this case the first link in the chain of equilibria: 

CuS (solid) ±5 CuS (dslvd) S=? Cu++ + S= U n . Mal ~ 
2HC1 £=» 2Cr + 2H+ j *- U2b {aslya) > 

tends so decidedly backward that only the use of concentrated acid 
will increase the concentration of the H + to an extent sufficient to 
secure even a slight advance of the whole action. We must add, 
therefore, to the above rule : provided also that the salt is not one 
of extreme insolubility. This point will be illustrated more fully in 
connection with the description of individual sulphides (see under 
Cadmium). 

Illustrations of the application of these generalizations are 
countless. Carbonic acid is made from marble (p. 381), hydrogen 
sulphide from ferrous sulphide (p. 272), hydrogen peroxide from 
sodium peroxide (p. 222), and phosphoric acid from calcium phos- 
phate (p. 370). In each case the acid employed to decompose the 
salt is more active than the acid to be liberated. On the other 



484 COLLEGE CHEMISTRY 

hand, calcium oxalate is insoluble in acetic acid because this acid 
is weaker than is oxalic acid. We have thus only to examine the 
list of acids showing their degrees of ionization (p. 241) in order 
to be able to tell which salts, if insoluble in water, will be dis- 
solved by acids and, in general, what acids will be sufficiently 
active in each case for the purpose. In chemical analysis we dis- 
criminate between salts soluble in water, those soluble in acetic 
acid (the insoluble carbonates and some sulphides, FeS and MnS, 
for example), those requiring active mineral acids for their solu- 
tion (calcium oxalate and the more insoluble sulphides, for ex- 
ample), and those insoluble in all acids (barium sulphate and 
other insoluble salts of active acids) . 

Precipitation of Insoluble Salts in Presence of Acids, — ■ 

The converse of solution, namely, precipitation, depends upon the 
same conditions : an insoluble salt which is dissolved by a given acid 
cannot be formed by precipitation in the presence of this acid. Thus, 
calcium oxalate can be precipitated in presence of acetic acid, but 
not in presence of active mineral acids in ordinary concentrations. 
Cupric sulphide or barium sulphate can be precipitated in pres- 
ence of any acid, but ferrous sulphide and calcium carbonate only 
in the absence of acids. 

From this it does not follow that calcium oxalate, for example, 
cannot be precipitated if once an active acid has been added to the 
mixture. To secure precipitation, all that is necessary is to re- 
move the excess of hydrogen-ion which is repressing the ionization 
of the oxalic acid. This can be done by adding a base, which re- 
moves the H + , or even by adding sodium acetate. The acetate- 
ion C 2 H 3 02~ unites with the H + to form the little ionized acetic 
acid, in presence of which calcium oxalate can be precipitated. 

Bleaching Powder Ca(OCl)Cl. — This substance (cf. p. 309) 
is manufactured by conducting chlorine into a box-like structure 
containing slaked lime spread upon perforated shelves. When the 
transformation has reached the limit (it is never complete), some 
lime dust is blown into the chamber to absorb the remainder of 
the free chlorine. 

That bleaching powder is a mixed salt CaCl(OCl) rather than an 
equimolar mixture of calcium chloride and calcium hypochlorite, 



CALCIUM 485 

which would have the same composition, CaCl 2 ,Ca(OCl) 2 , is 
proved by the facts that the material is not deliquescent as is 
calcium chloride, and that calcium chloride cannot be dissolved out 
of it by alcohol. 

Bleaching powder is somewhat soluble in water, and in solution 
the ions Ca++, Cl~ and CIO" are all present. Addition of active 
acids causes the formation of hydrochloric and hypochlorous acids 
(p. 309). Weak acids like carbonic acid displace the hypo- 
chlorous acid only (cf. p. 310), and hence the dry powder, when 
exposed to the air, has the odor of hypochlorous anhydride C1 2 
rather than that of chlorine. 

The substance is largely used by bleachers (cf. p. 311), and as a 
disinfectant to destroy germs of putrefaction and disease. 

Calcium Sulphate. — This salt is found in large quantities in 
nature. The mineral anhydrite CaSC>4 occurs in the salt layers. 
It contains no water of hydration, and its crystals belong to the 
rhombic system. The dihydrate, CaS0 4 ,2H 2 0, is more plentiful. 
In granular masses it constitutes alabaster. When perfectly crys- 
tallized (monoclinic system, Fig. 47, p. 83) it is named gypsum or 
selenite. The same hydrate is formed by precipitation from solu- 
tions. Its solubility is about 1 in 500 at 18°. 

Plaster of paris 2CaS04,H 2 is manufactured by heating gyp- 
sum until nearly all the water of hydration has been driven 
out. When it is mixed with water, the dihydrate is quickly re- 
formed and a rigid mass is produced. That the plaster sets 
rapidly, is due to the fact that the hemihydrate is more soluble 
than the dihydrate, and so a constant solution of the one and 
deposition of the other goes on until the hydration is complete. 
It becomes rigid, instead of forming a loose mass of dihydrate, 
because the process results in the formation of an interlaced and 
coherent mass of minute crystals. 

2CaS0 4 ,H 2 (solid) ^2CaS0 4 (dslvd) j ^ 2[CaS o 4 , 2 H 2 0] (solid). 

Plaster of paris is used for making casts and in surgery. The 
setting of the material is accompanied by a slight increase in 
volume, and hence a very sharp reproduction of all the details of 
the mold is obtained. An " ivory" surface, which makes washing 



486 COLLEGE CHEMISTRY 

practicable, is conferred by painting the cast with a solution of 
paraffin or stearin in petroleum ether. The waxy material, left 
by evaporation of the volatile hydrocarbons, fills the pores and 
prevents solution and disintegration of the substance by water. 
Stucco is made with plaster of paris and rubble, and is mixed with 
a solution of glue instead of water. 

Calcium Sulphide CaS. — This compound is most easily made 
by strongly heating pulverized calcium sulphate and charcoal. 
The sulphate is reduced: 4C + CaS0 4 -> CaS + 4CO. Calcium 
sulphide is meagerly soluble in water, but is nevertheless slowly 
dissolved in consequence of its decomposition by hydrolysis 
into calcium hydroxide and calcium hydrosulphide Ca(SH) 2 . 
It is used as a depilatory. Hair and wool are composed of proteins, 
which are decomposed by, and dissolved in alkaline solutions, 
like that here formed. Since calcium sulphide is thus decomposed 
by water it cannot be precipitated from aqueous solution by 
adding a soluble sulphide. 

Ordinary calcium sulphide, after it has been exposed to sunlight, 
usually shines in the dark. Barium sulphide behaves in the same 
way. On this account these substances are used in making 
luminous paint. They apparently owe this behavior to the 
presence of traces of compounds of vanadium and bismuth, for 
the purified substances are not affected in the same fashion. 

Phosphates of Calcium. — The tertiary orthophosphate of 
calcium Ca3(P0 4 )2, known as phosphorite, is found in many locali- 
ties, and is often derived from the remains of animals. Guano 
contains some of the same substance, along with nitrogen either 
in the form of organic compounds or as niter (p. 348). Apatite, 
3Ca 3 (P0 4 ) 2 ,CaF 2 , a double salt with calcium fluoride (or chloride), 
is a common mineral and frequent component of rocks. The 
orthophosphate forms about 83 per cent of bone ash, "and is 
contained also in the ashes of plants. It may be precipitated 
by adding a soluble phosphate to a solution of a salt of 
calcium. 

Since it is a salt of a weak acid, and belongs to the less insoluble 
class of such salts, calcium phosphate is dissolved by dilute mineral 
acids (c/. p. 483), the ions HP0 4 = and H 2 P0 4 ~ being formed. 



CALCIUM 487 

When a base, such as ammonium hydroxide, is added to the solu- 
tion, the calcium phosphate is reprecipitated (cf. p. 484). 

Calcium phosphate is chiefly used in the manufacture of phos- 
phorus and phosphoric acid (p. 370), and as a fertilizer. The 
supply of calcium phosphate in the soil arises from the decompo- 
sition of rocks containing phosphates, and is gradually exhausted 
by the removal of crops. Bone ash is sometimes used to make 
up the deficiency. It is almost insoluble in water, however, and, 
although somewhat less insoluble in natural water containing 
salts like sodium chloride, is brought into a condition for absorp- 
tion by the plants rather slowly. The "superphosphate" (see 
below) is much more soluble. 

Primary calcium orthophosphate ("superphosphate") is manu- 
factured in large quantities from phosphate rock by the action of 
sulphuric acid. The unconcentrated "chamber acid" is used for 
this purpose, as water is required in the resulting action. The 
amounts of material employed correspond to the equation: 

Ca 3 (P0 4 ) 2 + 2H 2 S0 4 + 6H 2 -* Ca(H 2 P0 4 ) 2 ,2H 2 + 2CaS0 4 ,2H 2 0. 

As soon as mixture has been effected, the action proceeds with 
evolution of heat, and a large cake of the two hydrated salts re- 
mains. This mixture, after being broken up, dried, and packed 
in bags, is sold as "superphosphate of lime." The primary phos- 
phate which it contains is soluble in water, and is therefore of 
great value as a fertilizer. 

Calcium Cyanamide. — Calcium carbide (p. 379), when 
strongly heated with nitrogen, gives a mixture of calcium cyan- 
amide and carbon: 

CaC 2 + N 2 ->CaCN 2 + C, 

which is sold as nitro-lime for use as a fertilizer. When treated 
with hot water, the cyanamide is hydrolyzed into calcium carbon- 
ate and ammonia: 

CaCN 2 + 3H 2 -> CaC0 3 + 2NH 3 . 

In the soil the decomposition may not be so simple, but combined 
nitrogen is furnished in a form that can be absorbed by plants. 



488 



COLLEGE CHEMISTEY 



At Odda (Norway) the carbide is pulverized and placed in a 
cylindrical furnace (Fig. 119) holding 300-450 kg. The heat (800- 
1000°) is supplied by the passage of an electric current through 
a thin carbon rod. The nitrogen" is obtained by 
the fractionation of liquid air and final removal 
of all oxygen by passage over heated copper, and 
is forced in under pressure. After thirty-six hours, 
nitrogen is no longer absorbed, and the charge is 
pulverized when cold. 

Sodium cyanide is now manufactured by fusing 
nitro-lime with sodium carbonate: 



CaCN 2 + C + Na 2 C0 3 -> CaC0 3 + 2NaNC. 



The cyanide is extracted from the insoluble cal- 
cium carbonate with water, in which it is exceed- 
ingly soluble. Sodium cyanide has now displaced potassium cyanide 
in the extraction of gold from its ores. 



Fig. 119. 



Nutrition and Fertilization of Crops. — A plant constructs 
its cellulose, starch, and sugar, and secures the carbon-part of 
all its organic contents from the carbon dioxide of the air (p. 387). 
The water (90-95 per cent of the total weight of the plant) comes 
from the soil and brings up in solution the other elements re- 
quired. All soils are able to supply sufficient magnesium, calcium 
and iron as bicarbonates. But the soil may lack: sulphur, 
absorbed as sulphates; nitrogen, absorbed chiefly as nitrates, but 
occasionally as salts of ammonium; potassium, as sulphate, 
chloride, or nitrate; and phosphorus, as soluble phosphates. The 
soil may be originally deficient in one or more of these necessary 
plant foods, or the supply may have been exhausted byrepeatec 
cropping. Every crop permanently removes certain quantities. 
For example, in the case of nitrogen, which is required to fori 
proteins that enter largely into the fruit (i.e., usually, the edible 
part), each crop of Indian corn (45 bushels) removes 63 pounc 
per acre, a crop of cabbage (15 tons) removes 100 pounds per acre, 
clover hay (2 tons) 82 pounds, and wheat (15 bushels) 31 pounds. 
When the store in the soil become meager, the crops become poor, 
and finally cost more for labor than they are worth. 



CALCIUM 489 

Thus, crops have to be fed, just like cattle. Moreover, the 
elements must be furnished in soluble form (cf. pp. 137, 340, 422). 
Fertilizers containing potassium (p. 445, 451) and phosphorus 
(p. 487) must be used, when the soil is deficient in these elements. 
The nitrogen fertilizers we have mentioned are sodium nitrate 
(p. 347), calcium nitrate (p. 353), ammonium sulphate (p. 411), 
guano and manure (p. 348), " tankage'' and ground bones from 
slaughter houses, calcium cyanamide (p. 487), and finally the 
nitrates from bacterial decomposition of root nodules (p. 339). 
That systematic use of fertilizers does influence the crops is indi- 
cated by the results of cultivation of land which, but for fertili- 
zation, would long since have become almost valueless. The 
wheat crop per acre, being the average of ten successive years is: 
Denmark 40 bushels, Great Britain 33, Germany 29, United 
States 14. 

Hard Water. — As we have seen (pp. 384, 476), limestone 
(solubility, 0.013 g. per liter), magnesium carbonate (sol'ty 1 g. 
per liter), and iron carbonate, although very insoluble, are acted 
upon by the carbonic acid in natural waters, giving bicarbonates 
which are roughly about thirty times as soluble. When the 
water is boiled, the actions are reversed, and the carbonates 
are reprecipitated. These bicarbonates constitute temporary 
hardness, and their decomposition produces "fur" in a kettle and 
boiler crust in a boiler. 

The sulphates of calcium (sol'ty 2 g. per liter) and of magnesium 
(sol'ty 354 g. per 1) are also commonly found in natural waters. 
These salts are not affected by mere boiling (as distinct from 
evaporation) and so, along with magnesium carbonate (1 g. per 1.) 
and calcium carbonate (0.013 g. per 1.) give permanent hardness 
to the water. 

Hardness is estimated in "degrees." In France, and com- 
monly in the laboratory, 1 part of CaC0 3 (or its equivalent of 
other salts) per 100,000 (0.01 g. per liter) constitutes one degree. 
In the United States one degree is 1 grain per gallon of 58,333 
grains (0.017 g. per 1.). In Britain one degree is 1 grain per gallon 
of 70,000 grains (0.014 g. per 1.). Well water, originating in 
chalk or limestone formations, may have 37° (Fr.) or more of 
hardness. 



490 COLLEGE CHEMISTRY 

Damage Due to Hardness in Water, — When hard water is 
continually fed into a steam boiler and only steam comes out, 
naturally the salts accumulate and produce in time a heavy 
boiler crust, which settles on the tubes. Being a poor conductor 
of heat compared with iron, this crust, if one-fourth of an inch 
thick, will increase the consumption (and cost) of fuel by 50 per 
cent. In addition, the iron, not being in direct contact with 
water, is heated to a higher temperature, and may even become 
red hot. It thus oxidizes more quickly on the outside, and dis- 
places hydrogen from water (or steam) on the inside (p. 52), 
thus changing on both sides gradually into the brittle magnetic 
oxide Fe 3 4 . If the crust is not removed, or prevented (see 
below), the life of the boiler is greatly shortened, and a serious 
explosion may even occur. 

In washing, in the household or laundry, much soap is wasted 
before the necessary lather is secured. The soap, for example, 
the sodium stearate (p. 415), gives magnesium and calcium stear- 
ates, which are insoluble, forming a curd: 

CaS0 4 + 2Na(C0 2 C 17 H 3 5) -> Ca(C0 2 C 17 H 3 5)2 I + Na 2 S0 4 . 

The permanent solution of soap, required for washing, does not 
begin to be formed until all the hardness has thus been precipi- 
tated. Hence, according to the equation, with 1° (U. S.) hard- 
ness, 100 gallons (U. S.) of water should use up 0.075 pound of 
soap (1° Brit, and 100 gal. Brit., 0.075 lb.). In point of fact, 
however, the colloidal calcium salts adsorb and carry down with 
them more than an equal amount of undecomposed soap. Hence, 
actual measurement shows that, with 1° (U. S. or Brit.) of hard- 
ness, 100 gallons (U. S. or Brit.) of water really destroy 0.17 
pound of soap. Thus, with 35°, no less than 6 pounds of soap 
per 100 gallons are wasted before the part of the soap that is to 
do the work begins to dissolve. 

Treatment of Hard Water. — The temporary hardness can 
be removed by boiling the water, or using some preheating 
arrangement in connection with the boiler (stationary engines 
only). 

Temporary hardness is commonly removed, on a large scale, by 
adding slaked lime (made into milk of lime) in exactly the quantity 



CALCIUM 491 

shown by an analysis of the water to be required, and stirring for 
a considerable time: 

Ca(HC0 3 ) 2 + Ca(OH) 2 -» 2CaC0 3 I + 2H 2 0. (1) 

The bicarbonate is neutralized and all the lime precipitated. 
The latter is removed by filtration. 

Permanent hardness is not affected by slaked lime, but is pre- 
cipitated by adding sodium carbonate in the necessary proportion: 

CaS0 4 + NaaCOs -* CaC0 3 J, + Na 2 S0 4 . (2) 

When both kinds of hardness are present, crude caustic soda 
(sodium hydr oxide) may be employed. It neutralizes the bicar- 
bonate, precipitating CaC0 3 : 

Ca(HC0 3 ) 2 + 2NaOH -> CaC0 3 j + Na 2 C0 3 + 2H 2 0. (3) 

and giving sodium carbonate. The latter then acts as in equa- 
tion (2). 

Instead of this, the treatments indicated in equations (1) and 
(2) may be applied in combination (Porter-Clark process).* 

In the new permutite process the water is simply filtered through 
an artificial sodium silico-aluminate (permutite NaP) which is 
supplied in the form of a coarse sand. The calcium, etc., in the 
water is exchanged for sodium, which does no harm: 

Ca(HC0 3 ) 2 + 2NaP -> 2NaHC0 3 + CaP 2 . 

After twelve hours' use, the permutite is covered with 10 per 
cent salt solution, and allowed to remain for the other twelve 
hours of the day, when it is ready for employment once more: 

2NaCl + CaP 2 -> CaCl 2 + 2NaP. 

Only salt, which is inexpensive, is consumed, and calcium chloride 
solution is thrown away. Permutite removes magnesium, iron, 
manganese, and other elements in the same way. The life of a 
charge is said to be over twenty years. 

Hard Water in the Laundry, — As we have seen (p. 490), 
soap will soften water, but the calcium and magnesium salts of the 
soap acids, which are precipitated, are sticky, and soil the goods 

* So far as the hardness is due to magnesium bicarbonate, a double propor- 
tion of lime must be added to precipitate the magnesium as hydroxide (sol'ty 
0.01 g. per 1.), because the carbonate is too soluble (1 g. per 1.). 



492 COLLEGE CHEMISTRY 

being washed. Other substances that soften water not only give 
non-adhesive precipitates, but are also much cheaper, and an at- 
tempt is generally made to utilize them. The use of slaked lime 
is impracticable on a small scale. 

Washing soda Na 2 CO 3 ,10H 2 O is added to precipitate both kinds 
of hardness: 

Ca(HC0 3 ) 2 + Na 2 C0 3 ->CaC0 3 + 2NaHC0 3 , 
CaS0 4 + Na 2 C0 3 -> CaC0 3 + Na 2 S0 4 . 

The small amounts of salts of sodium which remain in the water 
have no action on soap. 

Household Ammonia NH 4 OH acts like sodium hydroxide 
(p. 491): 

Ca(HC0 3 ) 2 + 2NH 4 OH -> CaC0 3 + (NHO2CO3 + 2H 2 0, 
CaS0 4 + (NH 4 ) 2 C0 3 ->CaC0 3 + (NH 4 ) 2 S0 4 . 

except that it will not precipitate magnesium-ion. 

Borax Na 2 B 4 O 6 ,10H 2 O (p. 432) is hydrolyzed and the sodium 
hydroxide contained in its solution acts as already (p. 491) de- 
scribed. 

The supposed bleaching or whitening action of borax or soda 
is a myth; these salts prevent staining by the iron in the water. 
They simply precipitate the iron, present as Fe(HC0 3 ) 2 , which 
almost all waters contain, as FeC0 3 , before the goods are put in. 
This precipitate is easily washed out in rinsing. The palmitate, 
etc., of iron, however, which the soap itself would throw down, is 
sticky and adheres to the cloth. The air subsequently oxidizes 
it (see p. 633) and gives hydrated ferric oxide (rust), which is 
brownish-red. 

It is evident that, properly to achieve their purpose, the soda 
and borax must be added, must be completely dissolved, and 
must be allowed to produce the precipitation of FeC0 3 , CaC0 3 , 
etc., all before the soap (or the goods) is introduced. If the soap 
is dissolved before or with the soda, it will take part in the pre- 
cipitation, and give sticky particles containing the iron and cal- 
cium salts of the soap acids. 

The soda, borax, and ammonia do not themselves remove dirt 
— that is done by the dissolved soap (p. 418). With the help of 
rubbing, however, they do emulsify and remove animal or vege- 



CALCIUM 493 

table oil and grease, but not mineral oil (p. 420), when these 
happen to be on the goods. But soap alone will do this also, and 
remove mineral oil as well. 

Washing powders are, or ought to be, mainly sodium carbonate, 
mixed with more or less pulverized soap. 

Calcium Silicate CaSi0 3 . — Calcium metasilicate CaSiOa 
forms the mineral wollastonite, which is rather scarce, but enters 
into the composition of many complex minerals, such as garnet 
and mica. It may be made by precipitation with a solution of 
sodium metasilicate (p. 428), or by fusing together powdered 
quartz and calcium carbonate: 

Si0 2 + CaC0 3 -> CaSiOs + C0 2 . 

Glass. — Common glass is a complex silicate of sodium and cal- 
cium, or a homogeneous mixture of the silicates of these metals 
with silica. It has a composition represented approximately by 
the formula Na 2 0,CaO,6Si0 2 , and is made by melting together 
sodium carbonate, limestone, and pure sand: 

Na^COa + CaC0 3 + 6Si0 2 -* Na 2 0,CaO,6Si0 2 + 2C0 2 . 

For the most fusible glass, a smaller proportion of sand is em- 
ployed. This variety is known, from its components, as soda- 
glass, or, from its easy fusibility, as soft glass. Plate-glass is made 
by casting the material in large sheets, rolling the sheet flat while 
hot, and polishing the surfaces when cold until they are plane. 
Window-glass is prepared by blowing bulbs of long cylindrical 
shape, and ripping them down one side with the help of a diamond. 
The resulting curved sheets are then placed on a flat surface in a 
furnace and are there allowed to open out. Beads are made, 
chiefly in Venice, by cutting narrow tubes into very short sections 
and rounding the sharp edges by fire. Ordinary apparatus is 
made of soft soda-glass, and hence when heated strongly it tends 
to soften and also to confer a strong yellow tint (cf. p. 465) on the 
flame. Bottles are made with impure materials, and owe their 
color chiefly to the silicate of iron which they contain. In all 
cases the articles are annealed by being passed slowly through a 
special furnace in which their temperature is lowered very grad- 



494 COLLEGE CHEMISTRY 

ually. Glass which has been suddenly chilled is in a state of 
tension and breaks easily when handled. 

Soft glass is partially dissolved by water. When powdered 
glass is shaken with water, sodium silicate dissolves at once, 
and in amount sufficient to give an alkaline reaction with phenol- 
phthalein (c/. p. 258). 

Bohemian, or hard glass, is much more difficult to fuse than 
soda-glass, and is also much less soluble in water. It is manu- 
factured by substituting potassium carbonate for the sodium 
carbonate. Specially insoluble glass, for laboratory use, such as 
Jena and non-sol glass, is made with boric anhydride B2O3, in ad- 
dition to silica, and some zinc oxide, so that it contains borates as 
well as silicates. When lead oxide is employed instead of lime- 
stone, a soda-lead glass known as flint glass is produced. This 
has a high specific gravity, and a great refracting power for light, 
and is employed for making glass ornaments. By the use of 
grinding machinery, cut glass is made from it. Engraving on 
glass is done with the sand blast. 

Colored glass is prepared by adding small amounts of various 
oxides to the usual materials. The oxides combine with the 
silica, and produce strongly colored silicates. Thus, cobalt oxide 
gives a blue, chromium oxide or cupric oxide a green, and uranium 
oxide a yellow glass. Cuprous oxide, with a reducing agent, and 
compounds of gold, give the free metals, suspended in colloidal 
solution (p. 416) in the glass, and confer a deep-red color upon it. 
Milk-glass contains finely powdered calcium phosphate in sus- 
pension, and white enamels are made by adding stannic oxide. 

Glass is a typical amorphous substance (pp. 266, 393). From 
the facts that it has no crystalline structure, and that it softens 
gradually when warmed, instead of showing a definite melting- 
point, it is regarded as a supercooled liquid of extreme viscosity. 
Most single silicates crystallize easily, and have definite freezing- 
(and melting-) points. Glass may be regarded as a solution of 
several silicates. When kept for a considerable length of time at 
a temperature insufficient to render it perfectly fluid, some of its 
components crystallize out, the glass becomes opaque, and "de- 
vitrification" is said to have occurred. The word "crystal" 
popularly applied to glass is thus definitely misleading. 

So-called quartz-glass is made of fused silica (p. 428). 



STRONTIUM 495 

Calcium-ion Co,**: Analytical Reactions. — Ionic calcium 
is colorless. It is bivalent, and combines with negative ions. 
Many of the resulting salts are more or less insoluble in water. 
Upon the insolubility of the carbonate, phosphate, and oxalate are 
based tests for calcium-ion in qualitative analysis (see p. 538). 
The presence of the element is most easily recognized by the 
brick-red color its compounds confer on the Bunsen flame, and 
by two bands — a red and a green one — which are shown by 
the spectroscope. 

Strontium Sr 

The compounds of strontium resemble closely those of calcium, 
both in physical properties and in chemical behavior. 

Occurrence. — The carbonate of strontium SrC0 3 is found as 
strontianite (Strontian, a village in Argyleshire). The sulphate, 
celestite SrS04, is more plentiful. The metal may be isolated by 
electrolysis of the molten chloride. 

Compounds of Strontium. — The compounds are all made 
from the natural carbonate or sulphate. The former may be dis- 
solved directly in acids, and the latter is first reduced by means 
of carbon to the sulphide, and then treated with acids. 

Strontium chloride SrCl 2 ,6H 2 0, made in one of the above ways, 
is deposited from solution as the hexahydrate. The nitrate 
Sr(NOs)2 comes out of hot solutions in octahedrons which are 
anhydrous. From cold water the tetrahydrate Sr(N03) 2 ,4H 2 is 
obtained. The anhydrous nitrate is mixed with sulphur, charcoal, 
and potassium chlorate to make "red fire." The oxide SrO may 
be secured by igniting the carbonate, but it is obtained with 
greater difficulty than is calcium oxide from calcium carbonate. 
It is generally made by heating the nitrate or hydroxide. 

Strontium hydroxide Sr(OH) 2 is made by heating the carbonate 
in a current of superheated steam: 

SrC0 3 + H 2 -> Sr(OH) 2 + C0 2 . 

This action takes place more easily than does the mere disso- 
ciation of the carbonate, because the formation of the hydroxide 
liberates energy, and this partially compensates for the energy 



496 COLLEGE CHEMISTRY 

which has to be provided to decompose the carbonate. The 
lowering of the partial pressure of the carbon dioxide by the 
steam also contributes to the result (cf. pp. 476-477). A hydrate 
Sr(OH) 2 ,8H 2 crystallizes from water. 

Strontium-ion Sr++ is bivalent, and gives insoluble compounds 
with carbonate-ion, sulphate-ion, and oxalate-ion. The presence 
of strontium is recognized by the carmine-red color which its 
compounds give to the Bunsen flame (see also p. 498). Its 
spectrum shows several red bands and a very characteristic blue 
line. 

Barium Ba 

The physical and chemical properties of the compounds of 
barium recall those of strontium and calcium. All the com- 
pounds of barium which are soluble in water, or can be brought 
into solution by the weak acids of the digestive fluids, are poison- 
ous. 

Occurrence. — Like strontium, barium is found in the form of 
the carbonate, witherite BaCC>3, and the sulphate BaSCX, heavy 
spar or barite (Gk. fiapvs, heavy). The free metal, which is silver- 
white, may be obtained by electrolysis of the molten chloride. 

The compounds are made by treating the natural carbonate 
with acids directly, or by first reducing the sulphate with carbon 
to sulphide, or converting the carbonate into oxide, and then 
treating the products with acids. 

Compounds of Barium. — The precipitated form of barium 
carbonate BaC03 is made by adding sodium carbonate to the 
aqueous extract from crude barium sulphide (q.v.). Barium car- 
bonate demands so high a temperature (about 1500°) for the at- 
tainment of a sufficient dissociation tension, that special means 
is employed for its decomposition. It is heated with powdered 
charcoal (cf. p. 385) : 

BaC0 3 + C->BaO + 2CO. 

Natural barium sulphate BaS0 4 is the source of many of the 
compounds. The precipitated sulphate, made by adding sul- 
phuric acid to the aqueous extract from barium sulphide, is used in 
making white paint (" permanent white"); in filling paper for 



BARIUM 497 

glazed cards, and sometimes as an adulterant of white lead. A 
mixture of barium sulphate and zinc sulphide ZnS, prepared in a 
special way, is called lithopone: 

BaS + ZnS0 4 -» BaS0 4 j + ZnS j . 

Made into paint it has greater covering power than white lead, 
does not darken with hydrogen sulphide as does the latter, and is 
non-poisonous. Barium sulphate is highly insoluble in water and 
is hardly at all affected by aqueous solutions of any chemical 
agents. 

Barium sulphide BaS, like the sulphides of calcium and stron- 
tium (p. 273), is very slightly soluble in water, but slowly passes 
into solution owing to hydrolysis and formation of the hydroxide 
and hydrosulphide. It is made by heating the pulverized sul- 
phate with charcoal: BaS0 4 + 4C -> BaS + 4CO. 

Barium chloride BaCl2,2H 2 is manufactured by heating the 
sulphide with calcium chloride. The whole treatment of the 
heavy spar is carried out in one operation: 

BaS0 4 + 4C + CaCl 2 -> 4CO + BaCl 2 + CaS. 

By rapid extraction with water, the chloride can be separated 
from the calcium sulphide before much decomposition of the 
latter (cf. p. 461) has taken place. 

Barium chlorate Ba(C10 3 )2 is made by treating the precipitated 
barium carbonate with a solution of chloric acid. It is deposited 
in beautiful monoclinic crystals, and is used with sulphur and 
charcoal in the preparation of " green fire." 

Barium monoxide BaO is manufactured from the carbonate (see 
above) or, in pure form, by heating the nitrate. The oxide unites 
vigorously with water to form the hydroxide. 

The monoxide, when heated in a stream of air or oxygen, gives 
barium peroxide: 2BaO + 2 <=* 2Ba0 2 , as a compact gray mass. 
This change and its reversal constitute the basis of Brin's process 
for obtaining oxygen from the air. At a suitable, high temper- 
ature, the air is forced in under pressure, causing the action to go 
forward, while the nitrogen escapes by a valve at the far end of 
the apparatus. Then, without change of temperature, by re- 
versing the pumps, oxygen is taken out, and the reaction goes 
backwards. This alternation makes the process a continuous 



498 COLLEGE CHEMISTRY 

one. A hydrate, Ba0 2 ,8H 2 0, is thrown down as a crystalline 
precipitate when hydrogen peroxide solution is added to a solu- 
tion of barium hydroxide: 

Ba(OH) 2 + H 2 2 ±± Ba0 2 j + 2H 2 0. 

Barium peroxide is used in the manufacture of hydrogen peroxide. 

Barium hydroxide Ba(OH) 2 , is made by union of the oxide with 
water, or by leading moist carbon dioxide over the sulphide and 
decomposing the resulting carbonate with superheated steam 
(p. 495). It is the most soluble of the hydroxides of this group, 
and gives, therefore, the highest concentration of hydroxide- 
ion. The solution is known as "baryta-water." It is also the 
most stable of the three hydroxides, and may be melted with- 
out decomposition. A hydrate Ba(OH) 2 ,8H 2 crystallizes from 
water. 

Barium nitrate Ba(NOs) 2 is made by the action of nitric acid on 
the sulphide, oxide, hydroxide, or carbonate of barium. The 
crystals from aqueous solution are anhydrous. 

Analytical Reactions of the Calcium Family. — Barium-ion 
Ba ++ is a colorless, bivalent ion. Many of its compounds are in- 
soluble in water, and the sulphate is insoluble in acids also. The 
spectrum given by the salts contains a number of green and orange 
lines. 

In solutions of salts of calcium, strontium, and barium, the ions 
may be distinguished by the fact that calcium sulphate solution 
will precipitate the strontium and barium as sulphates, but will 
leave salts of calcium in dilute solution unaffected. Similarly, 
strontium sulphate solution precipitates barium sulphate, and 
does not give any result with salts of the first two. The oxalate 
of calcium is precipitated in presence of acetic acid, while the 
oxalates of strontium and barium are not (cf. p. 484), and there 
are other differences of a like nature in the solubilities of the salts. 

Exercises. — 1. Arrange the chromates of the metals of this 
family in the order of solubility (see Table). Compare the solubili- 
ties with those of the carbonates, oxalates, and sulphates of the 
metals of the same family. 

2. What is meant by fluorescence (cf. any book on physics)? 



BARIUM 499 

3. What will be the ratio by volume, at 150°, of the nitrogen 
peroxide and oxygen given off by the decomposition of calcium 
nitrate? What would be the nature of the difference between the 
ratio at 150° and that at room temperature (c/. p. 352)? 

4. Apply the rule of precipitation to the case of adding sodium 
carbonate to a solution of barium chloride. 

5. Explain in terms of the ionic hypothesis the precipitation of 
the sulphate of strontium by calcium sulphate solution, and the 
absence of precipitation when the latter is added to a dilute solu- 
tion of a soluble salt of calcium. 

6. What inference do you draw from the fact that the oxalates 
of barium and strontium are not precipitated in presence of acetic 
acid, while the oxalate of calcium is so precipitated? Is the infer- 
ence confirmed by reference to the solubility data? 

7. Explain the fact that strontium and calcium chromates are 
easily dissolved by acetic acid, while barium chromate is dissolved 
only by active mineral acids. 

8. Explain the fact that all the carbonates, save those of 
sodium, potassium, and thallium, are precipitated in neutral solu- 
tions, but not in acidified solutions. Why is the precipitation 
incomplete when carbon dioxide is led through solutions of salts 
of the metals, but more complete when the hydroxides of the 
metals are used? 

9. Construct a table for the purpose of comparing the proper- 
ties of the free elements of this family and also the properties of 
their corresponding compounds. 

10. Are the elements of this family typical metals (p. 436)? 



CHAPTER XXXVII 

COPPER, SILVER, GOLD 

The three metals of this family, being found free in nature, are 
amongst those which were known in early times. They are the 
metals universally used for coinage and for ornamental purposes. 
They are the three best conductors of electricity (p. 436). 

The Chemical Relations of the Copper Family. — Copper 

(Cu, at. wt. 63.6), silver (Ag, at. wt. 107.88), and gold (Au, at. wt. 
197.2) occupy the right side in the second column of the table of 
the periodic system, and the chemical relations of these elements 
are in many ways in sharp contrast to those of the alkali metals, 
their neighbors, on the left side: 

Alkali Metals Coppek, Silver, Gold 

Very active; rapidly oxidized by air; Amongst least active metals; only 
displace all other metals from com- copper is oxidized by air; displaced 

bination (E.-M. series, p. 60). by most other metals. 

All univalent and give but one series Cu 1 and Cu 11 : two series. Ag 1 : one 
of compounds. Halides all soluble series. Au* and Au 111 : two series, 

in water. Chlorides of univalent series insol. 

Oxides and hydroxides strongly basic, Oxides and hydroxides feebly basic 
and halides not hydrolyzed (p. 437). (except Ag 2 0) ; halides hydrolyzed 

(except Ag-halides). Hence, basic 
salts are numerous. 
Never found in anion. Give no com- Frequently in anion, e.g., K.Cu(CN) 2 , 
plex cations. KAg(CN) 2 , KAu0 2 , K.Au(CN) 2 , 

and also in complex cations, e.g., 
Ag(NH 3 ) 2 .OH and Cu(NH 3 ) 4 .(0H) 2 

On account of their inactivity towards oxygen, and their easy 
recovery from combination by means of heat, silver and gold, 
together with the platinum family, are known as the "noble 
metals. " 

500 



COPPER 501 

Copper Cu 

Chemical Relations of the Element. — Copper is the first 
metallic element showing two valences which we have encountered. 
In such cases two more or less complete, independent series of salts 
are known. These are here distinguished as cuprous (univalent) 
and cupric (bivalent) salts. The methods by which a compound of 
one series may be converted into the corresponding compound of 
the other series should be noted. 

The chief cuprous compounds are Cu 2 0, CuCl, CuBr, Cul, 
CuCN, Cu 2 S. The cuprous compound is in each of these cases 
more stable (p. 93) than the corresponding cupric compound, and 
is formed from the latter either by spontaneous decomposition, 
as in the cases of the iodide and cyanide (2CuI 2 — > 2CuI + I2), or 
on heating. The cuprous halides and cyanide are colorless and 
insoluble in water. Cuprous-ion Cu + seems to be colorless. The 
cuprous salts of oxygen acids have few applications. 

The cupric compounds are more numerous, as they include also 
stable and familiar salts of oxygen acids, like CUSO4, Cu(NOs) 2 , etc. 
The anhydrous salts are usually colorless or yellow, but cupric-ion 
CU++ is blue, and so, therefore, are the aqueous solutions of the 
salts. The cupric are more familiar than the cuprous compounds, 
since cupric oxide, sulphate, and acetate are the compounds of 
copper which most frequently find employment in chemistry and 
in the arts. All the soluble salts of copper are poisonous. 

In addition to (1) having two valences Cu 1 and Cu 11 , and there- 
fore two series of compounds (two oxides, two chlorides, etc.), each 
of these states of copper also joins with other elements to form (2) 
complex positive ions such as Cu(NH 3 ) 2 + and Cu(NH 3 )4 ++ , just as 
hydrogen and ammonia form the complex positive ion NH4+, and 
the univalent form also gives (3) stable complex negative ions such 
as Cu(CN) 2 ~, CuCl 2 ~. None of the metallic elements discussed in 
the two preceding chapters showed any of these peculiarities. 
Many of the metals to be discussed later exhibit one or more of 
them, however. Especial attention should therefore be given to the 
chemistry of copper, in order that the behavior which such relations 
entail may be mastered at the first encounter, and the same rela- 
tions may be instantly recognized and understood when they 
reappear in other connections. 



502 COLLEGE CHEMISTRY 

There is only one other peculiarity which a metallic element fre- 
quently shows, although copper does not exhibit it. This is (4) the 
ability of its hydroxide to be, not only basic, as metallic hydroxides 
by definition (p. 436) must be, but also acidic. This behavior we 
encounter first in the case of gold (see p. 520) and in simpler and 
more familiar form in the case of zinc (see next chapter). 

Occurrence. — Copper is found free in the Lake Superior 
region. The sulphides, copper pyrites CuFeS 2 and chalcocite 
Cu 2 S, are worked in Montana, Utah, southwest England, Spain, 
and Germany. Malachite, Cu 2 (OH) 2 C0 3 (= Cu(OH) 2 ,CuC0 3 ), a 
basic carbonate, is mined in Arizona, Siberia, and elsewhere. 
Cuprite or ruby copper Cu 2 is also an important ore. 

Extraction from Ores. — For isolating native copper it is 
only necessary to separate the metal, by grinding and washing, 
from the rock through which it ramifies, and to melt the almost 
pure powder of copper with a flux (p. 438). The carbonate and 
oxide ores require coal, in addition, for the removal of the oxygen. 

The liberation of copper from the sulphide ores is difficult, and 
often involves very elaborate schemes of treatment. This arises 
from the fact that many copper ores contain a large amount of the 
sulphides of iron, and these have to be removed by conversion into 
oxide (by roasting) and then into silicate (with sand) . The silicate 
forms a flux, and separates itself from the molten mixture of copper 
and copper sulphide. In Montana it is found possible to abbrevi- 
ate the treatment. The ore is first roasted until partially oxidized. 
It is then melted in a cupola or a reverberatory furnace, and placed 
in large iron vessels like Bessemer converters (q.v.) provided with a 
lining rich in silica. A blast of air mixed with sand is next blown 
through the mass. The iron is completely oxidized to FeO and 
made into silicate FeSi0 3 , the sulphur escapes as sulphur dioxide, 
and arsenic and lead are likewise removed by this treatment. The 
silicate of iron floats as a slag upon the copper when the contents of 
the converter are poured out. The resulting copper is pure enough 
to be cast in large plates and purified by electrolysis (see p. 511). 

Much copper ore is of low grade, containing perhaps only 2 per 
cent of copper ore and 98 per cent of rock material. From such 
ores the usual methods of washing often recover only 70 per cent 



COPPER 503 

or less of the copper ore present, and 30 per cent or more is lost. 
The froth flotation process raises the proportion recovered to 85 or 
90 per cent of the whole. The finely crushed ore is agitated with 
water, to which is added some cheap oil and sometimes a little 
sulphuric acid. The mixture is then allowed to flow into a larger 
tank of water, in which the rock material immediately sinks to the 
bottom while the particles of ore are contained in the oily froth 
which rises to the top. The plant also occupies less than one-tenth 
of the space, and uses less than half the power required for treating 
the same amount of ore by washing. 

The world's production (1913) is about a million metric tons, of 
which the United States furnished 58 per cent, South America 11, 
Japan 6, and Germany 4. 

Physical and Chemical Properties. — Copper is red by re- 
flected and greenish by transmitted light. It melts at 1083°, and 
therefore much more easily than pure iron (1530°). Sp. gr. 8.93. 

In ordinary air copper becomes slowly covered with a green basic 
carbonate (not verdigris, q.v.). It does not decompose water at 
any temperature or displace hydrogen from dilute acids (p. 60). 
The metal attacks oxygen acids (pp. 275, 354), however. Sea- 
water and air slowly corrode the copper sheathings of ships, giving 
the basic chloride Cu4(OH) 6 Cl 2 ,H 2 0(= 3Cu(OH) 2 ,CuCl 2 ,H 2 0), 
which is found in nature as atakamite. 

On account of its resistance to the action of acids, copper is used 
for many kinds of vessels, for covering roofs and ships' bottoms, 
and for coins. It furnishes also electrotype reproductions of 
medals, of engraved plates, of type, etc. (see p. 510). Great quan- 
tities of the metal are used in electrical plants and appliances. 

Alloys. — The qualities of copper are modified for special pur- 
poses by alloying it with other metals. Brass contains 18^0 per 
cent of zinc, and melts at a lower temperature (p. 134) than does 
copper. A variety with little zinc is beaten into thin sheets, giving 
Dutch metal ("gold leaf"). Bronze contains 3-8 per cent of tin, 
11 or more per cent of zinc, and some lead. It was used for making 
weapons and tools before means of hardening iron were known, and 
later, on account of its fusibility, continued to be employed for 
castings until displaced largely by cast iron (discovered in the eight- 



504 COLLEGE CHEMISTRY 

eenth century). Gun metal contains 10 per cent, and bell metal 
20-24 per cent of tin. German silver contains 19-44 per cent of 
zinc and 6-22 per cent of nickel, and shows none of the color of 
copper. In many of these alloys the metals are partly in the form 
of chemical compounds, such as Cu 3 Sn and Cu 2 Zn 3 . 

Cupric Chloride CuCl 29 2H 2 Q» — This compound is made by 
union of copper and chlorine, or by treating the hydrate or car- 
bonate [ with hydrochloric acid. The blue crystals of a hydrate, 
CuCl 2 ,2H 2 0, are deposited by the solution. The anhydrous salt 
is yellow. Dilute solutions are blue, the color ,of cupric-ion, but 
concentrated solutions are green on account of the presence of the 
yellow molecules (p. 249) . The aqueous solution is acid in reaction 
(p. 437). When excess of ammonium hydroxide is added to the 
solution, the basic chloride, cupric oxychloride Cu4(OH) 6 Cl 2 (see 
above), which is at first precipitated, redissolves, and a deep-blue 
solution is obtained (see p. 506). This on evaporation yields deep- 
blue crystals of hydrated ammonio-cupric chloride Cu(NH 3 ) 4 .Cl 2 , 
H 2 0. The deep-blue color of the solution, which is given by all 
cupric salts, is that of ammonio-cupric-ion Cu(NH 3 )4 ++ . The dry 
salt also absorbs ammonia, giving CuCl 2 ,6NH 3 , but a reduction of 
pressure results finally in the loss of all the ammonia. 

Cuprous Chloride CuCL — It may be made by boiling cupric 
chloride solution with hydrochloric acid and copper turnings: 

CuCl 2 + Cu->2CuCl, or Cu++ + Cu -> 2Cu+. 

The solution contains compounds of cuprous chloride with hydrogen 
chloride HCl,CuCl, or HCuCl 2 and H 2 CuCl 3 , which are decomposed 
when the acid solution is diluted with water. The cuprous chlo- 
ride is insoluble in water, and forms a white crystalline precipitate. 

The foregoing action is an illustration of the fifth kind of ionic 
chemical change, namely, that in which a change in valence (and also 
in the amount of the electrical charge), occurs, without any altera- 
tion in the composition of the ionic substance. For other illustra- 
tions see pp. 158 (Mn++++ + 4C1" -> Mn++ + 2C1~ + Cl 2 ), 501. 

Cuprous chloride is hydrolyzed quickly by hot water, giving, 
finally, red, hydrated cuprous oxide, Cu 2 0. When dry it is not 
affected by light, but in the moist state becomes violet and, finally, 



COPPER 505 

nearly black. The action is said to be 2CuCl — > CuCl 2 + Cu. In 
moist air it turns green, and is oxidized to cupric oxy chloride (p. 
504). It is dissolved by hydrochloric acid, giving the colorless 
complex acids HCuCl 2 and H 2 CuCl3 just mentioned (see below). 
The solution is oxidized by the air, turning first brown and then 
green, and finally depositing the cupric oxy chloride. It also has 
the power of absorbing carbon monoxide, to form a compound 
said to be Cu(CO)Cl,H 2 0, and the property is used to separate 
this gas in analyzing mixtures of gases. Cuprous chloride also 
dissolves (see p. 506) in ammonium hydroxide, giving ammonio 
cuprous chloride Cu(NH 3 ) 2 .Cl, the ion Cu(NH 3 ) 2 + being colorless. 
The solution is quickly oxidized by the air, turns deep-blue, and 
then contains Cu(NH 3 )4 ++ . 

The Bromides and Iodides of Copper. — By treatment of 
copper with bromine-water, and slow evaporation of the solution, 
jet-black crystals of anhydrous cupric bromide CuBr 2 are obtained 
(cf. p. 249). When dry cupric bromide is heated, bromine is given 
off, and cuprous bromide CuBr remains. 

Cupric iodide Cul 2 appears to be unstable at ordinary tempera- 
tures. When a soluble iodide is added to a cupric salt, a white 
precipitate of cuprous iodide Cul and free iodine are obtained: 
2Cu++ + 4r^2CuIj + I 2 . 

The Solution of Insoluble Salts when Complex Ions are 
Formed. — The solution of an insoluble salt like cuprous chloride 
by hydrochloric acid or ammonium hydroxide is typical of a great 
variety of actions of which we here meet with the first examples. 
Compound or complex ions are formed (cf. p. 501). The explana- 
tion involves only principles already used in other cases. 

The dissolving of cuprous chloride in hydrochloric acid (p. 505), 
to form soluble, complex, highly ionized acids like H.CuCl 2 is a 
typical case. The complex negative ion CuCl 2 ~ which is formed 
is very little dissociated (CuCl 2 ~?^ Cu++ 2C1~), and gives a 
smaller concentration of Cu + than does the insoluble cuprous chlo- 
ride. The ion-product of cuprous chloride, and the concentra- 
tion relations of the ionic substance CuCl 2 ~ and its dissociation 
products (Cu + and 2C1~) are symbolized as follows: 

[cu*] x [en = r [Cu ^[ f = k. 



506 COLLEGE CHEMISTRY 

The value of [Cu + ] from cuprous chloride (first formula) is, in 
general, greater than its value from the ion CuCl 2 ~ of HCuCl 2 
(second formula), when excess of HC1 is present. Hence, the Cu + 
tends to pass over into the more stable compound, where it is more 
completely combined. More CuCl dissolves to replace the Cu+ 
which has been removed, and the change stops when the CuCl is 
all dissolved, or the values of [Cu + ] from both compounds have 
become equal. Thus, the complex ion is formed at the expense of 
the Cu + of the insoluble cuprous chloride, and the latter goes into 
solution progressively in the effort to restore the balance : 

CuCl (solid) ±=> CuCl (dsivd)^cr +cu+ l->cun-(<Mvd) 

2HC1 ^2H+ + 2Cr}^- UlU2 (dslvcU * 

The same exact laws of equilibrium used in discussing the dissolv- 
ing of salts by acids (p. 481) may be applied to the whole procedure. 

The dissolving of cuprous chloride by the free ammonia of ammo- 
nium hydroxide is explained in the same way. The only difference 
is that here the copper is in a complex positive ion. The ion 
Cu(NH3) 2 + gives little Cu + — less than does cuprous chloride, in 
spite of the insolubility of the latter. Hence the salt passes into 
solution until the ion-product [Cu + ] X [Cl~], with continually in- 
creasing [Cl~], reaches the value for a saturated solution or until 
the solid is exhausted. 

The deep-colored ion Cu(NH 3 )4 ++ given by cupric chloride and 
other cupric salts is also very little ionized. Hence ammonium 
hydroxide dissolves all the insoluble cupric compounds save only 
cupric sulphide, which is the most insoluble of all — that is, the one 
giving the smallest concentration of cupric-ion. Conversely, the 
sulphide is the only insoluble compound of copper which can be 
precipitated from ammoniacal solution. 

Foregoing Explanation Restated. — We may restate the 
explanation by answering a question: Why does cuprous chloride 
interact with, and go into solution in hydrochloric acid? Because 
it forms a complex compound HCuCl2, and, with the concentrations 
usually employed, the molecular concentration of cuprous-ion in the 
solubility product of cuprous chloride is greater than the molecular 
concentration of the same ion in the solution of the complex com- 
pound. 



COPPER 507 

The answer in other cases takes the same form. Thus, for 
cupric hydroxide Cu(OH) 2 dissolving in ammonium hydroxide 
solution, substitute cupric hydroxide for cuprous chloride and 
Cu(NH 3 ) 4 (OH) 2 for HCuCl 2 . 

Cuprous Oxide Cu 2 0, — This oxide is red in color, and natural 
specimens show octahedral forms. It is produced by oxidation of 
finely divided copper at a gentle heat, or by the addition of bases to 
cuprous chloride, and is best made by the action of glucose (p. 404) 
on cupric hydroxide (see Fehling's solution, below). The simple 
hydroxide, CuOH, is unknown, but the above mentioned pre- 
cipitate is a hydrated oxide 4Cu 2 0,H 2 0, and yields Cu 2 when 
heated. 

Cuprous oxide is acted upon by hydrochloric acid, giving cu- 
prous chloride, or rather HCuCl 2 . It also dissolves in ammonium 
hydroxide, giving, probably, Cu(NH 3 ) 2 .OH, which is colorless. 

Cupric Oxide and Hydroxide. — Cupric oxide CuO (black) is 
formed by heating copper in a stream of oxygen or, in less pure 
form, by igniting the nitrate, carbonate, or hydroxide. When 
heated strongly it loses some oxygen, and is partly reduced to 
Cu 2 0. 

Cupric hydroxide Cu(OH) 2 is precipitated as a gelatinous sub- 
stance by addition of sodium or potassium hydroxide to a solution 
of a cupric salt: CU++ + 20H~ -> Cu(OH) 2 , When the mixture 
is boiled, the hydroxide loses water and forms a black hydrated 
cupric oxide Cu(OH) 2 ,2CuO(?). 

The hydroxide interacts with ammonium hydroxide, forming the 
soluble compound Cu(NH 3 )4.(OH) 2 , which has a deep-blue color. 
Cellulose (cotton or paper) is soluble in this solution, and is re- 
precipitated by sulphuric acid. Artificial silk is made by pressing 
the solution through dies into the precipitant. Paper and cotton 
goods, when passed first through one and then the other of these 
liquids, receives a tough, waterproof surface. 

Cupric hydroxide interacts with a solution of Rochelle salt, 
giving a soluble compound. The liquid is known as Fehling's 
solution, and is used in testing for, and estimating quantities of 
glucose (p. 404), and other reducing substances. Cuprous oxide 
is precipitated (see above). 



508 COLLEGE CHEMISTRY 

Cupric Nitrate Cu(N0 3 ) 29 6H 2 0. — The nitrate is made by- 
treating cupric oxide or copper with nitric acid (p. 354), and is 
obtained from the solution as a deliquescent, crystalline hexahy- 
drate. When dehydrated at 65° the salt is partly hydrolyzed, 
and a basic nitrate Cu 4 (OH) 6 (N0 3 ) 2 remains. 

Carbonate of Copper. — No normal carbonate (CuC0 3 ) can be 
obtained. A basic carbonate (malachite) is found in nature, and is 
precipitated by adding soluble carbonates to cupric salts: 

2CuS0 4 + 2Na 2 C0 3 + H 2 -» Cu 2 (OH) 2 C0 3 + 2Na 2 S0 4 + C0 2 . 
The carbonate, if formed, would be hydrolyzed by water (p. 437). 

Cyanides of Copper. — With potassium cyanide and a solution 
of a cupric salt, cupric cyanide Cu(NC) 2 is precipitated. This is 
not stable, however, and gives off cyanogen, leaving cuprous 
cyanide CuNC: 

2Cu(NC) 2 -> 2CuNC + C 2 N 2 . 

Cuprous cyanide is insoluble in water, but interacts with an excess 
of potassium cyanide solution, producing a colorless liquid, from 
which K.Cu(CN) 2 (= KCN,CuCN), potassium cuprocyanide, may 
be obtained in colorless crystals. The complex anion Cu(CN) 2 ~ is 
so little ionized to Cu + and 2CN~ that all insoluble copper com- 
pounds, including cupric sulphide, are dissolved by potassium 
cyanide; and none of them can be precipitated from the solution. 
Zinc is actually unable to displace copper from such a solution. 
The cause of the solution of the salts is the same as when the com- 
plex ions Cu(NH 3 ) 2 +, Cu(NH 3 ) 4 ++, and CuCJ 2 ~are formed (p. 505). 

Cupric Acetate. — By the oxidation of plates of copper, sepa- 
rated by cloths saturated with acetic acid (vinegar), a basic acetate 
of copper (verdigris) is obtained: 

6Cu + 8HC 2 H 3 2 + 30 2 -+2Cu 3 (OH) 2 (C 2 H 3 2 ) 4 + 2H 2 0. 

It is used in manufacturing green paint, is insoluble in water, and is 
unaffected by light. It dissolves in acetic acid, and green crystals 
of the normal acetate Cu(C 2 H 3 2 )2,H 2 are obtained from the 
solution. The basic acetate is used in preparing Paris green. 



COPPER 509 

Cupric Sulphate CuSO*. — This salt is obtained by heating 
copper in a furnace with sulphur, and admitting air to oxidize the 
cuprous sulphide. The mixture of cupric sulphate and cupric oxide 
which is formed is treated with sulphuric acid. The salt is also 
made by allowing dilute sulphuric acid to trickle over granulated 
copper, while air has free access to the material: 

2Cu + 2H2SO4 + 2 -> 2CuS0 4 + 2H 2 0. 

This is an example of the use of two reagents which separately have 
little or no action (cf. pp. 426, 431). When concentrated and very 
hot, sulphuric acid will itself act as the oxidizing agent (cf. p. 276). 

Cupric sulphate crystallizes as pentahydrate CuS04,5H 2 in blue 
asymmetric crystals (Fig. 52, p. 95), and in this form is called blue- 
stone or blue vitriol. The aqueous solution has an acid reaction 
(p. 437). The anhydrous salt is white, and can be crystallized in 
thin needles from solution in hot, concentrated sulphuric acid (cf. 
p. 95). Cupric sulphate is employed in copper-plating (see p. 
510), in batteries, and as a mordant in dyeing (q.v.). A minute 
proportion is added to drinking water, to destroy algce, which other- 
wise propagate in the reservoirs and give a disagreeable taste and 
odor to the water. A solution, mixed with milk of lime (Cu(OH) 2 
is precipitated), Bordeaux mixture, is largely used as a spray on 
grape vines and other plants to prevent the growttf of fungi. 

When ammonium hydroxide is added to cupric sulphate solu- 
tion, a pale-green basic sulphate Cu4(OH) 6 S0 4 (?) is first precipi- 
tated. With excess of the hydroxide, the blue Cu(NH 3 ) 4 ++ ion 
(p. 506) is formed, and crystals of ammonio-cupric sulphate 
Cu(NH 3 )4.S0 4 ,H 2 can be obtained from the solution. 

Cupric sulphate also combines with potassium and ammonium 
sulphates, giving double salts of the form CuS0 4 ,K 2 S0 4 ,6H 2 0, which 
are deposited in large, monosymmetric crystals from the mixed 
solutions. Double salts (p. 245) exist as such in the solid form 
only and, in water, are resolved into the components and their 
ions. 

The Sulphides of Copper. — Cuprous sulphide Cu 2 S occurs in 
nature in rhombic crystals of a gray, metallic appearance. It is 
the sulphide formed by direct union of the elements. Cupric 
sulphide CuS is deposited as a black precipitate when hydrogen 



510 



COLLEGE CHEMISTRY 



sulphide is led through a solution of a cupric salt. When heated, 
it leaves cuprous sulphide and sulphur vapor is given off. 

Analytical Reactions of Compounds of Copper, — The ion 

of ordinary cupric salts, cupric-ion Cu" 1-1 ", is blue, and that of 
cuprous salts, cuprous-ion Cu + , is colorless. Cuprous solutions, 
however, are easily oxidized by the air and become blue. In solu- 
tions containing cupric-ion, hydrogen sulphide precipitates cupric 
sulphide, even in presence of acids (p. 483). Bases throw down 
the blue hydroxide, and carbonates precipitate a green basic salt 
(p. 508). Potassium ferrocyanide gives the brown, gelatinous 
cupric ferrocyanide: 

2Cu.S0 4 + K4.Fe(CN) 6 ^±Cu 2 .Fe(CN) 6 J + 2K 2 S0 4 . 

A characteristic test is the formation of the deep-blue Cu(NH 3 )4 ++ 
ion with excess of ammonium hydroxide. This solution itself 
responds to certain precipitants only (e.g., H 2 S). Solutions of 
complex cuprous cyanides such as K.Cu(CN) 2 are colorless, and do 
not respond to any of the above tests. With microcosmic salt or 
borax (pp. 371, 433), copper compounds form a bead which is 
blue in the oxidizing part of the flame and becomes red and opaque 
(liberation of copper) in the reducing flame. 

Electrotyping. — When plates of platinum, connected with a 
battery, are immersed in cupric sulphate solution, copper is 

deposited on the cathode (negative 
pole). The sulphate-ion S04 = mi- 
grates (p. 229) towards the anode 
(positive pole) and there liberates 
sulphuric acid and oxygen (p. 228). 
If, however, the anode is made of 
copper, the S04 = migrates, but is not 
discharged. Instead, copper goes into 
solution (Fig. 120) as Cu++, in amount 
equal to that deposited on the other 
pole. Thus, the only changes are, (1) an increase in concen- 
tration of cupric sulphate round the positive pole anode, and 
(2) a transfer of copper from the copper anode to the cathode 
(see. p. 511). 




COPPER 



511 



A copper electrotype of a medal (or a page of type) is made by 
first preparing a cast of the medal in plaster of Paris, gutta percha, 
or wax. The surface of the cast is then rubbed with graphite, to 
render it a conductor, and the cast is then used as the cathode in a 
cell with a copper anode, like that just described. The deposit of 
copper, when heavy enough, is stripped off. In making book 
plates, the cast is made with wax, and the copper electrotype is 
strengthened and thickened by filling the back with melted lead.* 



TpazszaaaamBEBazBiaa 



Copper Refining. — The tenacity, ductility, and conductivity 
of copper are seriously affected by small amounts of impurities, 
such as cuprous oxide or sulphide, which are soluble in the molten 
metal. Arsenic amounting to 0.03 per cent lowers the conductance 
about 14 per cent. There are also silver and gold in smelter 
copper. Hence, a large proportion of the copper on the market is 
purified by electrolysis. The principle is the same as that used 
in electrotyping. Thin sheets of copper form the cathodes, and 
thick plates of copper the anodes. These 
are suspended alternately and close to- 
gether in large troughs, lined with lead, 
and filled with cupric sulphate solution 
(Fig. 121, diagrammatic, view from above). 
The cathodes are all connected with the 
negative wire of the dynamo, and the 
anodes with the positive one. The Cu* -1- 
is attracted to the cathodes and is deposited 
upon them. The SO4— migrates towards 
the anodes, where copper from the thick 
plate becomes ionized in equivalent 
amount. The stock of cupric sulphate 
thus remains the same, and the liquid is 
stirred to keep the sulphate from accumu- 
lating close to the anodes. The practical effect of the electrolysis is 
to carry copper across from one plate to the other. The cathodes 
are removed from time to time, and the deposit of copper is 
stripped from their surface. Fresh anodes are substituted when 
the old ones are eaten away. Since there is no final decomposi- 

* For newspapers, a plate is made from the cast of the type more quickly 
by means of melted stereotype metal (lead, antimony, tin; 82 : 15 : 3). 



uu/n/wui/i 



Fig. 121. 



512 COLLEGE CHEMISTRY 

tion of any cupric sulphate, the only electrical energy required 
is that necessary to overcome the friction of the moving ions. 
Hence, a very small difference in potential (less than 0.5 volts) 
is sufficient (see p. 549). 

The less active metals which are mixed with the copper in the 
anode are not ionized, because there is plenty of the more active 
copper to carry the current. These metals, and traces of sulphides, 
therefore, fall to the bottom of the vat as a sludge. Zinc and other 
metals more active than copper, however, are ionized. Conversely, 
at the cathode, the copper, being the least active metal present 
in ionic form, is alone deposited. There is no tendency to dis- 
charge zinc or hydrogen, for example, so long as there are plenty 
of the more easily discharged copper ions available (see. p. 549). 
In this way, copper, 99.8 per cent pure, is obtained, gold and silver 
are recovered from the sludge, and the bath liquid is removed 
from time to time for purification from the more active metals it 
acquires. 

Silver 

Chemical Relations of the Element. — This element presents 
a curious assortment of chemical properties. It differs from copper 
in having a strongly basic oxide, and in giving salts with active 
acids which are not hydrolyzed by water. In these respects it 
approaches the metals of the alkalies and alkaline earths. It 
resembles copper in entering into complex compounds, and in 
giving insoluble halides. It differs from both copper and the 
metals of the alkalies, and resembles gold and platinum, in that 
its oxide is easily decomposed by heat, with formation of the- free 
metal, and in the low position it occupies in the electromotive 
series and the consequent slight chemical activity of the free 
metal. 

Occurrence. — Native silver, usually scattered through a rocky 
matrix, contains varying amounts of gold and copper. Native 
copper always contains dissolved silver. Sulphide of silver (Ag2S) 
occurs alone and dissolved in galenite (PbS) . Smaller amounts of 
the metal are obtained from pyrargyrite Ag 3 SbS3, proustite 
Ag 3 AsS 3 , and horn-silver AgCl. The chief supplies come from 
California, Australia, and Mexico. 



SILVER 513 

Metallurgy. — The silver contained free, or as sulphide, in 
ores of copper and lead, is found in the free state dissolved in the 
metals extracted from these ores, and is secured by refining them. 
In the electrolytic refining of copper, silver is obtained from the 
mud deposited in the baths (p. 512). The proportion present in 
lead is usually small. Parke's process, by which the silver is 
separated from the lead, takes advantage of the fact that molten 
zinc and lead are practically insoluble in one another, while silver 
is much more soluble in zinc than in lead. Lead dissolves 1.6 per 
cent of zinc, and zinc 1.2 per cent of lead. The principle is the 
same as in the removal of iodine from water by ether (p. 129). 
The lead is melted and thoroughly mixed by machinery with a 
small proportion of zinc. After a short time the zinc floats to the 
top, carrying with it almost all of the silver, and solidifies at a 
temperature at which the lead is still molten. The zinc-silver 
alloy, largely a compound Ag 2 Zn 5 , is skimmed off, and heated 
moderately in a furnace to permit the adhering lead to drain way. 
The zinc is finally distilled off in clay retorts, and the lead remain- 
ing with the silver is removed by cupellation. This operation 
consists in heating the molten metal strongly in a blast of air. The 
lead is converted into litharge (PbO), which flows in molten con- 
dition over the edge of the cupel, and the silver is then cast. 

Ores of silver which do not contain much or any lead are often 
smelted with lead ores, and the product is treated as described 
above, but many other processes are in use. The gold, which goes 
with the silver in Parke's process, is separated electrolytically 
(p. 511). Plates of the silver-gold alloy form the anode, and silver 
nitrate solution the vat-liquid. The silver, being the more active 
metal, is ionized and deposited on the cathode, while the gold 
collects as a powder in a bag surrounding the anode. 

During the first half of the nineteenth century the world's total 
output of silver averaged only 643 tons per year. Up to 1870 a 
gram of gold could buy 15.5 g. of silver. Now that the produc- 
tion has reached 7800 tons, the same amount of gold purchases 
about 40 g. The chief sources (1911) are Mexico 2460 tons, United 
States 1880, Canada 1018, Europe 525. 

Physical Properties. — Pure silver is almost perfectly white. 
It melts at 960°. Its sp. gr. is 10.5. Its ductility is such that wires 



514 COLLEGE CHEMISTRY 

can be drawn so fine that 2 kilometers weigh only about 1 g. In 
the molten condition it absorbs mechanically about twenty-two 
times its own volume of oxygen, but gives up almost all of this as 
it solidifies. Fantastically irregular masses result from the 
"sprouting" or " spitting" which accompanies the escape of the 
gas. 

By addition of ferrous citrate to silver nitrate, a red solution and 
lilac precipitate of free silver can be made. The latter, after 
washing with ammonium nitrate solution, gives a red, colloidal 
solution in water (cf. p. 416). It is a negatively charged colloid, 
and is coagulated by bivalent positive ions. 

Silver is alloyed with copper to render it harder. The silver 
coinage of the United States and the continent of Europe has a 
"fineness of 900" (900 parts of silver in 1000), and that of Great 
Britain 925. Silver ornaments have a fineness of 800 or more. 

Chemical Properties. — Silver, when cold, is oxidized by 
ozone, but not by oxygen (see silver oxide) . It does not ordinarily 
displace hydrogen from aqueous solutions of acids. Sulphur com- 
pounds in the air tarnish the surface, producing Ag2S, as do also 
eggs, secretions from the skin (proteins, p. 422), and vulcanized 
rubber. Silver interacts with cold nitric acid and with hot, con- 
centrated sulphuric acid, giving the nitrate or sulphate of silver 
and oxides of nitrogen or of sulphur (pp. 354, 276). 

The Halides of Silver. — The chloride AgCl, bromide AgBr, 
and iodide Agl are formed as curdy precipitates when a salt of silver 
is added to a solution containing the appropriate halide ion. The 
first is white, and melts at about 457°. The second and third are 
very pale-yellow and yellow respectively. The insolubility in 
water, which is very great, increases in the above order. 

When exposed to light, the chloride becomes first violet (col- 
loidal silver, dispersed in the AgCl) and finally brown, chlorine 
being liberated. The bromide and iodide behave similarly. Solid 
silver chloride absorbs ammonia, forming at low pressures 2AgCl,- 
3NH 3 , and with higher pressures of ammonia AgCl,3NH 3 . 

Complex Compounds of Silver. — Silver chloride dissolves 
easily in excess of ammonium hydroxide, giving the complex cation 



SILVER 515 

Ag(NH 3 )2 + - The bromide, which is less readily soluble, gives the 
same complexion. The iodide is hardly soluble at all. Ammonio- 
argentic-ion Ag(NH 3 ) 2 + , in solutions of concentrations such as are 
commonly used (O.liV to N), gives about the same concentration of 
argention Ag + as does the bromide, and much more than the highly 
insoluble iodide (cf. p. 506). Hence the latter is almost insoluble 
in ammonium hydroxide, and can be precipitated in ammoniacal 
solution. All three of the insoluble halides interact with solutions 
of potassium cyanide and of sodium thiosulphate, and go into 
solution, as do also all the other insoluble silver salts. Usually an 
equivalent amount of the cyanide or thiosulphate suffices, but for 
complete interaction with the sulphide an excess is required. With 
the cyanide, double decomposition gives first the insoluble silver 
cyanide AgCN, which then dissolves, forming the soluble potassium 
argenticyanide K.Ag(CN) 2 . The thiosulphate gives a solution 
containing the complex salt Na 3 .Ag(S 2 3 )2. The more active 
metals, like zinc and copper, displace silver from all solutions, 
whether the solutions contain simple or complex salts. 

Oxides of Silver. — When sodium hydroxide is added to a 
solution of a salt of silver, a pale-brown precipitate is obtained, 
which, after being freed from water, is found to be argentic oxide 
Ag20, and not AgOH. The aqueous solution of argentic oxide, 
however, is distinctly alkaline, and presumably therefore does 
contain the hydroxide: 2 AgOH <=± Ag 2 + H 2 0. It is an active 
basic oxide. When moist, it absorbs carbon dioxide from the air. 
With ammonium hydroxide it forms the soluble Ag(NH 3 ) 2 .OH. 
When the oxide is heated, it gives off oxygen, leaving metallic silver. 
The action is reversible and at 302° the dissociation pressure of the 
oxygen is 20.5 atmospheres. At a higher pressure than this, there- 
fore, oxygen will combine with silver (at 302°). 

Silver peroxide Ag20 2 is formed by the action of ozone on silver. 
In the electrolysis of silver nitrate a deposit of shining black 
crystals which contain some silver peroxide is formed on the 
anode. 

Salts of Silver, — Silver nitrate AgN0 3 is obtained by treating 
silver with aqueous nitric acid: 

3Ag + 4HN0 3 -> 3AgN0 3 + NO + 2H 2 0. 



516 COLLEGE CHEMISTRY 

From the solution, colorless rhombic crystals are deposited. 
These melt at 208.6°. Thin sticks made by casting (lunar caustic) 
are used to cauterize sores, because the substance combines with 
proteins to form insoluble compounds. The aqueous solution is 
neutral. The pure salt is not affected by light, but when deposited 
on cloth, on the skin of the fingers, or on the mouth of the reagent 
bottle, it is reduced by organic matter, and silver is liberated. For 
this reason it is an ingredient in some marking-inks. 

Silver carbonate, the neutral salt Ag 2 C0 3 , and not a basic car- 
bonate, is precipitated from solutions of salts of silver by soluble 
carbonates. It is slightly yellow in color. With water it gives a 
faint alkaline reaction and, like calcium carbonate, is soluble ir 
excess of carbonic acid (p. 384). When heated, the carbonate 
decomposes, leaving metallic silver. The sulphate Ag 2 S0 4 is made 
by the action of concentrated sulphuric acid on the metal. When 
it is mixed with a solution of aluminium sulphate (q.v.) } octahedral 
crystals of silver-alum Ag 2 S0 4 , A1 2 (S0 4 ) 3 ,24H 2 are obtained. Silver 
sulphide Ag 2 S is precipitated by hydrogen sulphide from solutions 
of all silver compounds, whether free acids are present or not, and 
irrespective of the form in which the silver is combined. Excess of 
potassium cyanide, however, prevents its precipitation from the 
argenticyanide. The sulphide is formed by the action of metallic 
silver on alkaline hydrosulphides, and this interaction forms the 
basis of the "hepar" test for sulphur. Silver orthophosphate 
Ag 3 P0 4 (yellow), arsenate Ag 3 As0 4 (brown), and chromate Ag 2 Cr0 4 
(crimson) are produced by precipitation, and their distinctive 
colors enable us to use silver nitrate in analysis as a reagent for 
identifying the acid radicals. 

Electroplating. — The process is similar to the electrode- 
position of copper (p. 510). The article to be plated is cleaned 
with extreme care and attached to the negative wire. A plate of 
silver forms the positive electrode and, since simple salts of silver 
do not give coherent deposits, the bath is a solution of potassium 
argenticyanide. The potassium-ion K + migrates to the negative 
wire and, since potassium requires a much greater E.M.F. for its 
liberation than does silver, silver is there deposited from the trace 
of argentic-ion given by the complex silver ions in the neighbor- 
hood : Ag(CN) 2 " <=> Ag+ + 2CN", Ag+ + -> Ag°. 



SILVER 517 

At the positive electrode silver goes into solution in equivalent 
amount, giving argentic-ion, and the above equations are 
reversed. 

Mirrors are silvered through the reduction of ammonio-silver 
nitrate by organic compounds such as formaldehyde CH 2 (forma- 
lin), or grape sugar: 

4AgOH + CH 2 -> 3H 2 + 4Ag j + C0 2 . 
The film of silver is washed, dried, and varnished. 

Photography. — Bromo-gelatine dry plates are covered with 
an emulsion of gelatine in which silver bromide is suspended. 

After exposure, often for only a fraction of a second, there is no 
visible alteration in the film. The image is developed. Chemically, 
this consists in reducing the silver bromide to metallic silver by 
means of reducing agents. While the whole of the halide upon the 
plate is reducible, if the reducing agent is kept upon it for a suffi- 
cient length of time, the parts reached by the light are affected first, 
and with a speed proportional to the intensity of the illumination 
undergone by each part. The unreduced silver bromide is then 
dissolved out with sodium thiosulphate ("hypo"), and the silver 
image remains. It is also thus saved from being fogged over by the 
silver that would be deposited if the plate were to be brought into 
the light without this treatment (fixing). The result is a "nega- 
tive," as the parts brightest in the object are now opaque, and the 
darkest parts of the object are transparent. 

A common developer is the potassium salt of hydroquinone 
C 6 H4(OH) 2 , which gives quinone CeELiC^: 

2AgBr + (KO) 2 C 6 H4 -> 2Ag + 2KBr + Ce^O,. 

In printing, the fight and dark are again reversed, the denser 
parts of the negative protecting the compounds on the paper below 
it from action, and leaving them white. Either " bromide " papers 
(such as velox, invented by Baekeland), which require only brief 
exposure and are developed like the plate, are used, or silver 
chloride is the sensitive substance, and prolonged exposure to light 
is allowed to liberate the proper amount of silver. The operation 
of fixing is performed as before. In toning chloride papers, a solu- 



518 COLLEGE CHEMISTRY 

tion of sodium chloraurate is employed. A portion of the silver 
dissolves, displacing gold (p. 260), which is deposited in its place: 

NaAuCU + 3Ag -* NaCl + 3AgCl + Au. 
The thin film of gold gives a richer color to the print. 

Analytical Reactions of Silver Compounds. — Argentic-ion 
Ag + is colorless. Many of its compounds are insoluble, the pre- 
cipitation of the chloride, which is insoluble in dilute acids, being 
used as a test. Mercurous chloride and lead chloride are also white 
and insoluble, but silver chloride dissolves in ammonium hydroxide, 
mercurous chloride (q.v.) turns black, and lead chloride is not 
altered in color (and is also soluble in hot water). With excess 
of ammonium hydroxide, silver salts give the complex cation 
Ag(NH 3 ) 2 + and, from solutions containing silver in this form, only 
the iodide and sulphide can be precipitated. Sodium thiosulphate 
and potassium cyanide dissolve all silver salts, giving salts of 
complex acids with silver in the anion (p. 515). 

Gold Au 

Chemical Relations of the Element. — This element forms 
two very incomplete series of compounds corresponding respec- 
tively to aurous and auric oxides, Au 2 and Au 2 3 . The former is 
a feebly basic oxide, the latter mainly acid-forming. No simple 
salts with oxjrgen acids are stable. All the compounds of gold are 
easily decomposed by heat with liberation of the metal. All other 
common metals displace gold from solutions of its compounds (p. 
260). Mild reducing agents likewise liberate gold. The element 
enters into many complex anions. 

Occurrence and Metallurgy. — Gold is found chiefly in the 
free condition, disseminated in veins of quartz, or mixed with 
alluvial sand. Small quantities are found also in sulphide ores of 
iron, lead, and copper. Telluride of gold (sylvanite), in which 
silver takes the place of a part of the gold [Au,Ag]Te 2 , is found in 
Colorado. 

From the alluvial deposits, gold is usually separated by washing 
in a cradle (sp. gr., gold 19.32, rock about 2.6), as in the Klondyke. 



GOLD 519 

Quartz veins, which in the Transvaal Colony reach a thickness of 
a meter and carry an average of 18 g. of gold per ton, are mined, 
and the material is pulverized with stamping machinery. About 
55 per cent of the gold is then separated by allowing the powdered 
rock to be carried by a stream of water over copper plates amal- 
gamated with mercury. The gold dissolves in the latter, and is 
secured by removal and distillation of the amalgam. The 45 per 
cent of finer particles, contained in the sludge which runs off 
(" tailings"), are extracted by adding a dilute solution of sodium 
cyanide (MacArthur-Forest process) and exposing the mixture to 
the air. Oxidation and simultaneous interaction with the cyanide 
give sodium aurocyanide NaAu(CN) 2 . From this solution the 
gold is isolated, either by electrolysis, or in the form of a purple 
powder by precipitation with zinc. The same cyanide is used for 
another batch. 

The gold separated from ores in the above ways contains silver, 
copper, lead, and other metals, and various methods of refining, 
mainly electrolytic, are used. 

The world's production of gold during the first half of the nine- 
teenth century averaged 27 tons annually. In 1897 it was 363 
tons, and in 1899, 472.6 tons. It is partly this rapid increase in the 
supply of gold (which is our standard of value) which has made it 
relatively cheaper, and other articles more expensive. In 1913 
the total production was 680 tons, of which the Transvaal gave 
40 per cent, the United States 20 per cent, and Australia 12 per 
cent. 

Properties of the Metal. — Gold is yellow in color, and is the 
most malleable and ductile of all the metals. It melts at 1063°. 
Its sp. gr. is 19.32. To give it greater hardness it is alloyed with 
copper, the proportion of gold being defined in " carats." Pure 
gold is "24-carat." British sovereigns are 22-carat and contain 
■fe of copper. American, French, and German coins are 21.6-carat, 
or 90 per cent gold. 

Gold is not affected by free oxygen nor by hydrogen sulphide. 
It does not displace hydrogen from dilute acids, nor does it interact 
with nitric or sulphuric acids or any oxygen acids except selenic 
acid. It combines, however, with free chlorine and bromine. It 
interacts with a mixture of nitric and hydrochloric acids (aqua 



520 COLLEGE CHEMISTEY 

regia), giving chlorauric acid H.AuCl4(= HCl,AuCl 3 ). This 
happens, not because aqua regia is more active than are any of the 
substances it contains, but because it furnishes both the chlorine 
and the chloride-ion Cl~~ required to produce the exceedingly stable 
(little dissociated) anion AuCLT. Chlorine-water (C1 2 ,H+,C1~,- 
C10~) dissolves it also, for the same reason. Gold is the least 
active of the familiar metals. 

Compounds with the Halogens, — Chlorauric acid, formed 
as above, is deposited in yellow, deliquescent crystals of 
H.AuCl4,4H 2 0. The yellow sodium chloraurate NaAuCl4,2H 2 0, 
obtained by neutralization of the acid, is used in photography (p. 
518). The acid gives up hydrogen chloride when heated very 
gently, leaving the red, crystalline auric chloride AUCI3. When 
dissolved in water, this gives H 2 AuCl 3 0. When auric chloride is 
heated to 180°, aurous chloride AuCl and chlorine are formed. 
This salt is a white powder. It is insoluble in water, but in boiling 
water is converted quickly into auric chloride and free gold: 
3AuCl -> 2Au + AuCl 3 + H 2 -> H 2 AuCl 3 0. 

Other Compounds. — When caustic alkalies are added to 
chlorauric acid, or to sodium chloraurate, auric hydroxide Au(OH) 3 
is precipitated. This substance is an acid, and interacts with 
excess of the base, forming aurates. These are derived from met- 
auric acid (Au(OH) 3 — H 2 = HAu0 2 ),as, for example, potassium 
aurate K.Au0 2 ,3H 2 0. This salt interacts by double decomposition 
giving, for instance, with silver nitrate, the insoluble silver salt 
AgAu0 2 . Its solution is alkaline in reaction, showing that auric 
acid is a weak acid (cf. p. 437). Auric oxide Au 2 3 is a brown, and 
aurous oxide Au 2 is a violet powder. On account of its reducing 
action, hydrogen sulphide precipitates from chlorauric acid a dark- 
brown mixture containing much aurous sulphide Au 2 S and free 
sulphur, as well as some auric sulphide Au 2 S 3 . 

The aurocyanides like K.Au(CN) 2 (= KCN,AuCN), and the 
auricyanides, like K.Au(CN) 4 (= KCN,Au(CN) 3 , are formed by the 
action of potassium cyanide on aurous and auric compounds, 
respectively. They are colorless and soluble. Their solutions are 
used as baths, in conjunction with a gold anode, for electro- 
gilding. 



GOLD 521 

It will be seen that gold, although physically a metal, is chemi- 
cally on the whole a nonmetaUic element. 

Assaying. — In assaying, the material containing the gold is 
heated with borax and lead in a small crucible (cupel) of bone ash. 
The lead and copper are oxidized, and their oxides are absorbed by 
the cupel, leaving a drop of molten alloy of gold and silver. The 
cold button is flattened by hammering and rolling, and treated with 
nitric acid to remove the silver. The gold, which remains un- 
attacked, is washed, fused again, and weighed. The acid will not 
interact with the silver, and remove it completely, if the quantity 
of gold exceeds 25 per cent. When the proportion of gold is 
greater than this, a suitable amount of pure silver is fused with the 
alloy ("quartation"). 

Exercises. — 1. Write equations for the interactions, (a) of salt 
water and oxygen with copper (p. 503), (6) of ferrous oxide and 
sand (p. 502). 

2. Write the formulae of the basic chloride, nitrate, carbonate, 
and sulphate of copper as if these substances were composed of the 
normal salt, the oxide and water (p. 369). 

3. Can you develop any relation between the facts that solu- 
tions of cupric salts are acid in reaction and that they give basic 
carbonates by precipitation? 

4. Formulate the action of potassium cyanide in dissolving 
cupric hydroxide and cuprous sulphide, assuming that potassium 
cuprocyanide is formed. 

5. How should you set about making cupric orthophosphate, 
ammonium cuprocyanide, and lead cuprocyanide? 

6. Write the formulae of some of the double salts analogous to 
potassium-cupric sulphate (p. 509). 

7. What chemical reagents are present in a Bunsen flame? If 
borax beads were made in the oxidizing flame with cupric chloride, 
cuprous bromide, and cupric sulphate, severally, what actions 
would take place? 

8. Write the equations for the interaction of, (a) silver and 
concentrated sulphuric acid, (6) silver chloride and sodium car- 
bonate when heated strongly, (c) sodium thiosulphate and silver 
bromide. 



522 COLLEGE CHEMISTRY 

9. What reagents should you use to precipitate the phosphate, 
arsenate, and chromate of silver? 

10. Write the equations for the interactions of, (a) potassium 
hydroxide and auric hydroxide, (6) potassium cyanide and sodium 
chloraurate. 

11. In what respects are the elements of this family distinctly 
metallic, and in what respects are they allied to the non-metals 
(p. 436)? 

12. Collect all the evidence tending to show that the cuprous 
compounds are more stable than the cupric. 

13. Make a classified list of the methods by which cupric com- 
pounds are transformed into cuprous, and vice versa. 

14. Of which metals should it be possible to obtain colloidal 
suspensions in water, and of which not (p. 260)? Suggest some 
liquids in which you should expect to obtain colloidal suspensions 
of the alkali metals. 



CHAPTER XXXVIII 

GLUCINUM, MAGNESIUM, ZINC, CADMIUM, MERCURY. 

THE RECOGNITION OF CATIONS IN QUALITATIVE 

ANALYSIS 

The Chemical Relations of the Family. — The remaining 
elements of the third column of the periodic table, namely glucinum 
or beryllium (Gl, or Be, at. wt. 9.1), magnesium (Mg, at. wt. 24.32), 
zinc (Zn, at. wt. 65.4), cadmium (Cd, at wt. 112.4), and mercury 
(Hg, at. wt. 200.6), although all bivalent, do not form a coherent 
family. Glucinum and magnesium resemble zinc and cadmium, 
and differ from the calcium family, in that the sulphates are soluble, 
the hydroxides easily lose water leaving the oxides, and the metals 
are not rapidly rusted in the air and do not easily displace hydrogen 
from water. They resemble the calcium family, and differ from 
zinc and cadmium, in that the sulphides are hydrolyzed by water, 
the oxides are not reduced by heating with carbon, complex cations 
are not formed with ammonia, and the metals do not enter into 
complex anions. But glucinum differs from magnesium and 
resembles zinc in that its hydroxide is acidic as well as basic. This 
is not unnatural, since in the periodic system it lies between 
lithium, a metal, and boron, a non-metal. Mercury is the only 
member of the group that forms two series of compounds. These 
are derived from the oxides HgO and Hg 2 0. Mercury approaches 
the noble metals in the ease with which its oxide is decomposed by 
heating, and in the position of the free element in the electromotive 
series. 

The vapor densities of zinc, cadmium, and mercury show the 
vapors of these three metals to be monatomic. 

Glucinum Gl 

Glucinum (or beryllium) is bivalent in all its compounds. Its 
oxide and hydroxide are basic, and are also feebly acidic towards 
active bases (see Zinc hydroxide). The element derives its name 
from the sweet taste of its salts (Gk. yXvKvs, sweet). 

523 



524 COLLEGE CHEMISTRY 

Glucinum occurs in beryl, a metasilicate of glucinum and alumin- 
ium Gl 3 Al 2 (Si0 3 )6. Beryls, tinted green by the presence of a little 
silicate of chromium, are known as emeralds. The metal, obtained 
by electrolysis of the easily fusible double fluoride G1F 2 ,2KF, burns 
when heated in the air. It displaces hydrogen from dilute acids, 
and when heated, from caustic potash: G1+2KOH— >K 2 G10 2 +H 2 . 

Magnesium Mg 
Chemical Relations of the Element, — Magnesium is 
bivalent in all its compounds. The oxide and hydroxide are basic 
exclusively. The element does not enter into complex cations or 
anions. 

Occurrence. — Magnesium carbonate occurs alone as magne- 
site, and in a double salt with calcium carbonate MgC0 3 ,CaC0 3 as 
dolomite. The sulphate and chloride are found as hydrates and as 
constituents of double salts (see below) in the Stassfurt deposits. 
Olivine is the orthosilicate Mg 2 Si0 4 . Talc (soapstone) is an acid 
metasilicate H 2 Mg 3 (Si0 3 )4. Serpentine is a hydrated disilicate, 
[Mg,Fe] 3 ,Si 2 07,2H 2 0, as is also meerschaum. Asbestos is an an- 
hydrous silicate. The element derives its name from Magnesia, a 
town in Asia Minor. 

The Metal. — Magnesium is manufactured by electrolysis of 
dehydrated and fused carnallite MgCl 2 ,KCl,6H 2 0. The iron 
crucible in which the material is melted forms the cathode, and a 
rod of carbon the anode. The metal is silver-white, and when 
heated can be pressed into wire and rolled into ribbon (m.-p. 651°). 

Chemically the metal is less active than are the metals of the 
alkaline earths. It slowly becomes coated with a layer of the car- 
bonate. It displaces hydrogen slowly from boiling water and 
rapidly from cold, dilute acids. Magnesium burns in air with a 
white light. The ash contains the nitride Mg 3 N 2 , as well as the 
oxide. 

Powdered magnesium is used in pyrotechny and, with potassium 
chlorate (10 : 17), in making flashlight powder for use in photog- 
raphy. 

Magnesium Chloride MgCl 2 ,6H z O. — This highly deliques- 
cent salt occurs in salt deposits, alone, and as carnallite MgCl 2> - 



MAGNESIUM 525 

KC1,6H 2 0. The latter is an important source of potassium chlo- 
ride (p. 445), and almost 'all the magnesium chloride combined 
with it is thrown away. When the hexahydrate is heated, a part 
of the chloride is hydrolyzed, some magnesium oxide remaining, 
and some hydrogen chloride being given off. Sea-water cannot be 
used in ships' boilers because of the hydrochloric acid thus liberated 
by the action of the magnesium chloride which the water contains. 
Anhydrous magnesium chloride MgCl 2 is obtained by heating the 
double chloride MgCl 2 ,NH4Cl,6H 2 0, for this salt can be dehydrated 
without hydrolysis of the chloride. The ammonium chloride is 
volatilized (p. 453). 

The Oxide and Hydroxide, — Magnesium oxide MgO is made 
by heating the carbonate, and is known as calcined magnesia. It 
is a white, highly infusible powder, and is used for lining electric 
furnaces and making crucibles. It combines slowly with water to 
form the hydroxide Mg(OH) 2 . 

The hydroxide is found in nature as brucite. It is also precipi- 
tated from solutions of magnesium salts by alkalies. It is very 
slightly soluble in water. The solution has a faint alkaline reaction. 
When magnesium chloride is added to the moist hydroxide, a 
hydrated basic chloride, (Mg(OH) 2 ) x ,(MgCl 2 ) y ,(H 2 0) z , is formed. 
The mixture, to which sawdust is sometimes added, is used as a 
plaster-finish in building. 

Magnesium hydroxide is not precipitated by ammonium hy- 
droxide when ammonium salts are present also. The ammonium 
salts, being highly ionized and giving a high concentration of 
ammonium-ion NELf 1 ", repress the ionization of the feebly ionized 
ammonium hydroxide, and so reduce the concentration of hy- 
droxide-ion which it furnishes. With the ordinary concentration 
of Mg 4 " 4 ", therefore, the amount of hydroxide-ion existing in pres- 
ence of excess of a salt of ammonium is too small to bring the 
solubility product [Mg ++ ] X [OH - ] 2 up to the value required for 
precipitation (c/. p. 479). Conversely, magnesium hydroxide 
interacts with solutions of ammonium salts and passes into solution: 

Mg(OH) 2 (solid) <=>Mg(OH) 2 (dslvd) fc> Mg^+20H"L 9NTT on 
2NH4C1 fc* 2C1" +2NH4+ J ^ iN tUUti ' 

In presence of excess of ammonium chloride, the OH~ combines 
with NH4 + to form molecular ammonium hydroxide, and the 



526 COLLEGE CHEMISTRY 

equilibria in the upper line are displaced forwards to generate a 
further supply of the OH~. With sufficiently great concentration 
of the ammonium chloride, all the magnesium hydroxide may thus 
dissolve. The whole case is analogous to the interaction of acids 
with insoluble salts (p. 480). 

Other Salts of Magnesium. — The normal carbonate MgC0 3 
is found in nature. Only hydrated basic carbonates are formed by 
precipitation, and their composition varies with the conditions. 
The carbonate, manufactured in large amounts and sold as mag- 
nesia alba, is approximately Mg 4 (OH) 2 (C0 3 )3.3H 2 0. It is used in 
medicine and as a cosmetic. 

The common heptahydrate of magnesium sulphate MgS04,7H 2 
crystallizes from cold water in rhombic prisms, and is called Epsom 
salts. It is efflorescent. The monohydrate MgS04,H 2 0, which 
remains, and is found also in the salt layers as kieserite, has a very 
low aqueous tension, and is not rapidly dehydrated except above 
200°. Magnesium sulphate is used in the manufacture of sodium 
and potassium sulphates, and is employed also for " loading' ' 
cotton goods, and as a purgative. 

The sulphide MgS may be formed by heating the metal with 
sulphur. It is insoluble in water, but is decomposed and gives, 
finally, hydrogen sulphide and magnesium hydroxide: 

MgS + 2H 2 <=± Mg(0H)4 + H 2 S. 

The only phosphate of importance is ammonium-magnesium 
orthophosphate NH4MgP04,6H 2 0, which appears as a crystalline 
precipitate when sodium phosphate and ammonium hydroxide 
(and chloride, p. 525) are mixed with a solution of a magnesium 
salt. 

Analytical Reactions of Magnesium Compounds. — The 

magnesium ion is colorless and bivalent. Soluble carbonates pre- 
cipitate basic carbonates of magnesium, but not when ammonium 
salts are present. The latter limitation distinguishes compounds of 
magnesium from those of the calcium family. Sodium hydroxide 
precipitates the hydroxide of magnesium, except when salts of 
ammonium are present. The mixed phosphate of ammonium and 
magnesium, in presence of ammonium hydroxide, is the least 
soluble salt. 



zinc 527 

Zinc Zn 

Chemical Relations of the Element, — Zinc is bivalent in all 
its compounds. Of these there are two sets, — the more numerous 
and important one, in which zinc is the positive radical (Zn.S04, 
Zn.Cl 2 , etc.), and a less numerous set, the zincates, in which zinc 
is in the negative radical (Na2.Zn0 2 , etc.). Both sets of salts are 
hydrolyzed by water, as the hydroxide is feeble whether it is con- 
sidered as an acid or as a base. The element also enters into 
complex cations and anions. The salts are all poisonous. 

Occurrence and Extraction from the Ores. — The chief 
sources of zinc are calamine Zn 2 Si0 4 ,H 2 0, smithsonite ZnC0 3 , 
zinc-blende (Ger. blenden, to dazzle) or sphalerite ZnS, franklinite 
Zn(Fe0 2 ) 2 , and zincite ZnO. 

The ores are first concentrated, recently by froth flotation (p. 
503) . They are then converted into oxide — the carbonate by 
ignition, and the sulphide by roasting. The sulphur dioxide is used 
to make sulphuric acid. A mixture of the oxide with coal is then 
distilled in earthenware retorts at 1300-1400°, the zinc condensing 
in earthenware receivers, while carbon monoxide burns at a small 
opening: 

2ZnS + 30 2 -> 2ZnO + 2S0 2 , 
ZnO + C -> CO + Zn. 

At first zinc dust (a mixture of zinc and zinc oxide) collects in the 
receiver, and afterwards liquid zinc. The product, which is cast in 
blocks, is called spelter. 

Properties and Uses of the Metal. — Zinc is a bluish-white 
crystalline metal. When cold it is brittle, but at 120-150° it can 
be rolled into sheets between heated rollers and then retains its 
pliability when cold. At 200-300° the metal becomes once more 
brittle, at 419° it melts, and at 925° it boils. The vapor at 1740° 
is monatomic. 

The metal burns in air with a greenish flame, giving zinc oxide. 
In cold, moist air it is very slowly oxidized, and becomes covered 
with a firmly adhering, non-porous layer of basic carbonate which 
protects it from further action. The metal displaces hydrogen 
from dilute acids. Zinc also attacks boiling alkalies, giving the 
soluble zincate (see below) : 2KOH + Zn -> K 2 Zn0 2 + H 2 . 



528 COLLEGE CHEMISTRY 

Sheet zinc, in consequence of its lightness (sp. gr. 7), is used in 
preference to lead (sp. gr. 11.5) for roofs, gutters, and architectural 
ornaments. Galvanized iron is made by dipping sheet iron, 
cleaned with sulphuric acid or the sand blast, into molten zinc. 
The latter, being more active (p. 260), is rusted instead of the iron, 
but the rusting is very slight (see above). Objects of iron, cleaned 
and baked in zinc dust, also acquire a coating of zinc (sherardizing). 
Zinc is used also in batteries and for making alloys (p. 435). It 
mixes in all proportions with tin, copper, and antimony. 

Zinc Chloride ZnCl z . — This salt is usually manufactured by 
treating zinc with excess of hydrochloric acid, evaporating the 
solution to dryness, and fusing the residue. When hydrochloric 
acid is thus present, the chloride ZnCl 2 is obtained. Evaporation 
of the pure aqueous solution, which is acid in reaction, results in 
considerable hydrolysis and formation of much of the basic chloride 
Zn 2 OCl 2 : z n ci 2 + H 2 <=± HC1 + Zn(OH)Cl, (1) 

2Zn(OH)Cl -* Zn 2 OCl 2 + H 2 0. (2) 

The salt is used in solid form as a caustic and, by injection of a 
solution into wood {e.g., railway sleepers), as a poison to prevent 
the growth of organisms which promote decay. In both cases the 
salt combines with proteins, forming solid products. The hot 
solution also dissolves cellulose (cotton or paper). When the 
solution is pressed through an orifice into alcohol, the cellulose is 
precipitated in the form of a thread. By carbonizing such threads, 
carbon filaments for incandescent lamps are made. 

Zinc Oxide and Hydroxide and the Zincates. — The oxide 

ZnO is obtained as a white powder by burning zinc or by heating 
the precipitated basic carbonate. It turns yellow when heated, 
recovering its whiteness when cold, in the same way that mercuric 
oxide is brown whilst hot and bright red when cold. It is em- 
ployed in making a paint — zinc-white or Chinese white — which 
is not darkened by hydrogen sulphide. It is used also as a filler in 
making rubber automobile tires. 

The hydroxide Zn(OH) 2 appears as a white, flocculent solid when 
alkalies are added to solutions of zinc salts. It interacts as a basic 
hydroxide with acids, giving salts of zinc: 

Zn(OH) 2 + H 2 S0 4 $=> Zn.S0 4 + 2H 2 0. 



zinc 529 

It also interacts with excess of the alkali employed to precipitate 
it, giving a soluble zincate, such as potassium zincate K^ZnC^: 

H 2 Zn0 2 t + 2KOH *± K 2 .Zn0 2 + 2H 2 0. 

Zinc hydroxide is ionized both as an acid and as a base : 

2H+ + Zn0 2 = ^ Zn(OH) 2 (dslvd) *± Zn++ + 20H" 

it 

Zn(OH 2 ) (solid) 

Substances which are both bases and acids are called amphoteric. 
The ionization as an acid is less than that as a base, but both are 
small. Addition of an acid like sulphuric acid, however, furnishes 
hydrogen-ion; the hydroxyl ions combine with this to form water, 
and all the equilibria are displaced to the right. With a base, on 
the other hand, the hydrogen-ion is removed and the basic ioniza- 
tion simultaneously repressed, so that the equilibria are displaced 
to the left. 

Zinc hydroxide interacts with ammonium hydroxide, giving the 
soluble ammonio-zinc hydroxide Zn(NH 3 ) 4 .(OH) 2 . The case is like 
those of copper (p. 507) and silver hydroxides (p. 515). 

Compounds of zinc, when heated in the Bunsen flame with a salt 
of cobalt, give a zincate of cobalt (Rinmann's green) CoZn0 2 . 

Other Salts of Zinc, — The normal zinc carbonate ZnCOs may 
be precipitated by means of sodium bicarbonate, but normal car- 
bonate of sodium gives basic carbonates, such as Zn 2 (OH) 2 C03: 
2ZnS0 4 + 2Na2C0 3 + H 2 -> Zn 2 (OH) 2 C0 3 + 2Na*S0 4 + C0 2 T. 

Zinc sulphate ZnS0 4 is formed when zinc-blende is roasted. It 
gives rhombic crystals of the hydrate ZnS0 4 ,7H 2 0. This, and 
the corresponding compounds of magnesium MgS0 4 ,7H 2 0, of iron 
FeS0 4 ,7H 2 0, and of other bivalent metals are known as vitriols. 
The zinc salt is white vitriol. It is used in cotton-printing and as 
an eye-wash (f per cent solution) . The sulphate gives double salts, 
such as potassium-zinc sulphate ZnS0 4 ,K 2 S0 4 ,6H 2 (cf. p. 509). 

Zinc sulphide ZnS is more soluble in water than is sulphide of 
copper, and hence it interacts with excess of strong acids, and 
passes into solution. It is not soluble enough, however, to be 
much affected by weak acids like acetic acid (cf. p. 483). Zinc 
sulphide is thus capable of being precipitated when acetic acid is 



530 COLLEGE CHEMISTRY 

present, or when hydrogen sulphide is led into a solution of the 
acetate of zinc: 

Zn(C 2 H 3 2 )2 + H 2 S +± ZnS j + 2HC 2 H 3 2 . 

But when an active acid is present, or is formed, the sulphide is 
precipitated incompletely or not at all, the action being reversible: 

ZnS0 4 + H 2 S <=* ZnS + H 2 S0 4 . 

There are thus two ways of obtaining the sulphide by precipita- 
tion. A soluble sulphide causes it to be thrown down completely, 
because no acid is liberated in the action: 

ZnCl 2 + (NH4) 2 S ?± ZnS | + 2NH4C1. 

The other method is to add sodium acetate to the solution of the 
salt, and then lead in hydrogen sulphide. The acid, liberated by 
the action upon the salt, interacts with the sodium acetate, giving 
a neutral salt of sodium and acetic acid, and the zinc sulphide is 
not much affected by the latter (c/. p. 484). For uses, see lithopone 
(p. 497). 

Analytical Reactions of Zinc Salts. — Zinc sulphide is pre- 
cipitated by the addition of ammonium sulphide to solutions of 
zinc salts and of zincates. Sodium hydroxide gives the insoluble 
hydroxide, which, however, interacts with excess of the alkali, 
giving the soluble zincate of sodium. Compounds of zinc, when 
heated on charcoal with cobalt nitrate, give Rinmann's green 
(p. 529). 

Cadmium Cd 

Chemical Relations of the Element. — This element is biva- 
lent in all its compounds. Its oxide and hydroxide are basic exclu- 
sively, and the salts are not hydrolyzed by water. It enters into 
complex compounds having the ions Cd(NH 3 )4 ++ and Cd(CN) 4 = . 
Note its resemblances to, and differences from zinc. 

The Metal. — Aside from the rare mineral greenockite CdS, 
cadmium is found in small amounts (about 0.5 per cent), as car- 
bonate and sulphide, in the corresponding ores of zinc. During 
the reduction, being more volatile than zinc, it distils over first 
(b.-p. 778°). The metal is white, and is more malleable than zinc. 



CADMIUM 531 

It displaces hydrogen from dilute acids (cf. p. 260). It is used in 
making fusible alloys. 

Compounds of Cadmium. — The chloride CdCl 2 ,2H 2 is efflo- 
rescent and is not hydrolyzed during dehydration or in solution. 
Zinc chloride (p. 528) is deliquescent and is easily hydrolyzed. 

The hydroxide Cd(OH) 2 is made by precipitation (white), and 
interacts with acids (as a basic hydroxide), but not at all with 
bases. It dissolves in ammonium hydroxide, however, forming 
Cd(NH 3 )4.(OH) 2 . The oxide CdO is a brown powder, obtained by 
heating the hydroxide, carbonate, or nitrate, or by burning the 
metal. 

The sulphate crystallizes from solution as 3CdS0 4 ,8H 2 0. 
Soluble carbonates throw down the normal carbonate of cadmium 
CdC0 3 . 

Hydrogen sulphide precipitates the yellow sulphide CdS even 
from acid solutions of the salts. The substance is used as a pig- 
ment. The sulphide of cadmium, however, is less insoluble in 
water (cf. p. 483) than are the sulphides of copper and mercury, 
and is not completely precipitated from a strongly acid solution 
(e.g., HC1 > 0.3iV). " 

The Solubilities of the Sulphides of the Metals. — The 

reader will remember the order of solubility of the metallic sul- 
phides more easily if he notes that it is practically the same as the 
order of activity of the free metals (p. 260 or Appendix V). Thus, 
the sulphides down to that of aluminium are dissolved by water 
(K 2 S and Na 2 S) or are decomposed by water (BaS, SrS, CaS, MgS, 
A1 2 S 3 ). The hydroxides formed, being soluble (except Al(OH 3 )), 
the whole dissolves except in the case of A1 2 S 3 . Zinc sulphide is 
insoluble in water, but is soluble enough to interact with (and 
dissolve in) dilute acids, even a feeble one like acetic acid. Ferrous 
sulphide requires a dilute active acid; cadmium sulphide requires 
a higher concentration of an active acid, as do also CoS and NiS; 
cupric sulphide requires an oxidizing acid like hot nitric acid; and 
mercuric sulphide resists even this. 

Analytical Reactions of Cadmium Compounds. — The cad- 
mium ion Cd 4-1 " is bivalent and colorless. The yellow cadmium 



532 COLLEGE CHEMISTRY 

sulphide is precipitated by hydrogen sulphide, even from acid 
solutions of the salts. The white, insoluble hydroxide is not 
soluble in sodium hydroxide. 

Mercury Hg 

Chemical Relations of the Element. — Like copper, this ele- 
ment enters into two series of compounds, the mercurous Hg 1 and 
the mercuric Hg 11 . The mercurous halides, like the cuprous halides 
(and the argentic halides), are insoluble in water and are decom- 
posed by light. Both of the oxides, EfeO and HgO, are basic 
exclusively, but in a feeble degree. The hydroxides, like silver 
hydroxide, are not stable, and lose water, giving the oxides. The 
salts of both sets are markedly hydrolyzed by water, and basic 
salts are therefore common. No carbonate is known. Mercury 
enters into the anions of a number of complex salts, such as 
HgCLfS Hgl 4 ~, Hg(CN) 4 ~, etc. It forms a class of ammono-basic 
mercury compounds, like Hg^NH 2 Cl, all of which are insoluble. 

The mercury salts of volatile acids, like the corresponding salts of 
ammonium (p. 345), can all be volatilized completely. Mercury 
vapor and all mercury compounds are poisonous, the soluble ones 
more markedly so than the insoluble ones. 

Occurrence and Isolation of the Metal. — Mercury occurs 
native and to a larger extent as red, crystalline cinnabar, mercuric 
sulphide HgS. The chief mines are in Spain, Italy, Austria, and 
California. 

The liberation of the metal is easy, because, when roasted, the 
sulphide is decomposed, and the sulphur forms sulphur dioxide. 
The mercury does not unite with the oxygen, for the oxide decom- 
poses (p. 14) at 400-600°: 

HgS + 2 -+Hg + S0 2 . 

In some places the ore is spread on perforated brick shelves in a 
vertical furnace, and the gases pass through tortuous flues in which 
the vapor of the metal condenses. 

Physical Properties. — Mercury or quicksilver (N.L. hydrargy- 
rum, from Gk. v8up, water, and apyvpos, silver) is a silver-white 
liquid. At - 38.7° it freezes, and at 357° it boils. 



MERCURY 533 

On account of its high specific gravity (13.6, at 0°) and low vapor 
tension, the metal is employed for filling barometers. Its uniform 
expansion favors its use in thermometers. It forms amalgams 
with all the familiar metals, with the exception of iron and plati- 
num. The latter, however, is "wet" by it (cf. pp. 345, 519). 
Compounds, such as NaHg2, are often present in amalgams. 

Chemical Properties, — When kept at a temperature near to 
its boiling-point, mercury combines slowly with oxygen. Mercury 
does not displace hydrogen from dilute acids (p. 260), but with 
oxidizing acids like nitric acid and hot concentrated sulphuric acid, 
the nitrates and sulphate (mercuric) are formed. With excess of 
mercury, mercurous nitrate, and with excess of the hot acid, mer- 
curic nitrate, are produced. When mercury is divided into minute 
droplets, with relatively large surface, it is used hi medicine ("blue 
pills"), and shows an activity which is entirely wanting in larger 
masses. 

The Halides of Mercury. — Mercurous chloride HgCl (calomel) 
is obtained as a white powder by precipitation. It is made by 
subliming mercuric chloride with mercury: 

HgCla + Hg*±2HgCl, 

or more usually by subliming a mixture of mercuric sulphate, made 
as described above, with mercury and common salt. It is de- 
posited on the cool part of the vessel as a fibrous crystalline mass. 
Its vapor is composed entirely of mercury and mercuric chloride. 
It is slowly affected by light just as is silver chloride. Here, how- 
ever, the chlorine which is released combines with another molecule 
of the salt to form mercuric chloride. The substance is used in 
medicine on account of its tendency to stimulate all organs pro- 
ducing secretions. 

By direct union with chlorine, mercuric chloride HgCl2 (corrosive 
sublimate) is formed. It is usually manufactured by subliming 
mercuric sulphate with common salt, and crystallizes in white, 
rhombic prisms. It melts at 265° and boils at 307°. The solu- 
bility at 20° is 7.4 : 100 Aq. The aqueous solution is slightly acid 
in reaction. The salt is easily reduced to mercurous chloride. 
When excess of stannous chloride is added to the solution, the 



534 COLLEGE CHEMISTRY 

white precipitate of calomel, first formed, passes into a heavy 
gray precipitate of finely divided mercury: 

2HgCl 2 + SnCl 2 -> SnCU + 2HgCl, 
2HgCl + SnCl 2 -> SnCU + 2Hg. 

Corrosive sublimate, when taken internally, is extremely poison- 
ous. A very dilute solution (1 : 1000) is used in surgery to destroy 
lower organisms and thus prevent infection of wounds. Mercuric 
chloride acts also as a preservative of zoological materials, form- 
ing insoluble compounds with proteins, and preventing decay. 
For the same reason, albumin (white of an egg) is given as an 
antidote in cases of sublimate poisoning. 

Mercurous iodide Hgl is formed by rubbing iodine with excess of 
mercury. It also appears as a greenish-yellow precipitate when 
potassium iodide is added to a solution of a mercurous salt. It 
decomposes spontaneously into mercury and mercuric iodide: 

2HgI*±Hg + HgI 2 . 

Mercuric iodide Hgl 2 is obtained by direct union of mercury with 
excess of iodine, or by addition of potassium iodide to a solution of 
a mercuric salt. It is a scarlet powder, insoluble in water, but 
soluble in alcohol and ether. It interacts with excess of potassium 
iodide, forming the soluble, colorless potassium mercuri-iodide 
K 2 .HgI 4 with which many precipitants fail to give mercury com- 
pounds. 

The Oxides. — When bases (excepting ammonium hydroxide, 
see p. 535) are added to solutions of mercurous salts, the greenish- 
black mercurous oxide Hg 2 is thrown down. The hydroxide is 
doubtless formed transitorily and then loses water (c/. Silver oxide, 
p. 515). Under the influence of light or gentle heat (100°), this 
oxide resolves itself into mercuric oxide and mercury. 

Mercuric oxide HgO is formed as a red, crystalline powder, when 
mercury is heated in air near to 357°, but is usually made by 
decomposing the nitrate. Commercial specimens, incompletely 
decomposed, thus give some nitrogen tetroxide when heated. 
It is formed also as a yellow powder by adding bases (except- 
ing ammonium hydroxide, see p. 535) to solutions of mercuric 
salts. 



MERCURY 535 

Other Salts of Mercury. — Mercurous nitrate HgN03,H 2 is 
formed by the action of cold, diluted nitric acid upon excess of 
mercury. It is hydrolyzed, slowly by cold, and rapidly by warm 
water, giving an insoluble basic nitrate: 

2HgN0 3 + H 2 *± HN0 3 + Hg 2 (OH)N0 3 | . 

On this account a clear solution can be made only when some nitric 
acid is added. Free mercury is also kept in the solution to reduce 
mercuric nitrate, which is formed by atmospheric oxidation: 

Hg(N0 3 ) 2 + Hg->2HgN0 3 , or Hg++ + Hg -> 2Hg+. 

Mercuric nitrate Hg(N0 3 ) 2 ,8H 2 is produced by using excess of 
warm, concentrated nitric acid with mercury. The aqueous solu- 
tion is strongly acid, and deposits a yellowish, crystalline, basic 
nitrate Hg 3 (OH) 2 0(N0 3 ) 2 . The hydrolysis is reversed by adding 
nitric acid. 

Mercurous sulphide Hg 2 S is formed by precipitation from mer- 
curous salts, but decomposes into mercury and mercuric sulphide. 

Crystallized mercuric sulphide HgS occurs as cinnabar, and is 
red. When formed by precipitation with hydrogen sulphide, or by 
rubbing together mercury and sulphur, it is black and amorphous. 
By sublimation, in the course of which it dissociates and recom- 
bines, the black form gives the red, crystalline one. 

The black and the red varieties do not interact with concen- 
trated acids, or even with boiling nitric acid, which oxidizes most 
sulphides readily. They are, therefore, still less soluble than is 
cupric sulphide (pp. 483, 531). They are attacked, however, by 
aqua regia, because of the formation of the negative ion (see gold, 
p. 520) of a complex salt H 2 .HgCl4 (= 2HCl,HgCl 2 ). The red 
form of the sulphide is used in making paint (vermilion) . 

Mercuric fulminate Hg(ONC) 2 is obtained as a white precipitate 
when mercury is treated with nitric acid, and alcohol is added to 
the solution. It decomposes suddenly when struck, and is used in 
making percussion caps and detonators. 

Ammono- Compounds of Mercury. — When ammonium 
hydroxide is added to a solution of a mercuric salt, a white sub- 
stance, of a type which we have not previously encountered, is 
thrown down. Mercuric chloride gives Hg(NH 2 )Cl, commonly 



536 COLLEGE CHEMISTRY 

called "infusible white precipitate," or ammono-basic mercuric 
chloride. 

HgCl 2 + H.NH 2 + NH 3 -» Hg(NH 2 )Cl + NH4CL 

The action is similar to an hydrolysis which gives a basic salt: 
HgCl 2 + H.OH->Hg(OH)Cl + HCl, excepting that ammonia 
H.NH 2 plays the part of the water. Water gives aquo-basic 
salts. When liquid ammonia is the solvent, ammono-basic salts 
are produced. In a few cases, as here, an ammono-basic salt is 
obtained even when water is present. The study of reactions in 
liquid ammonia solutions by E. C. Franklin has led to the discovery 
of a large number of new and most interesting substances. 

Mercuric nitrate Hg(N0 3 ) 2 and ammonium hydroxide give an 
insoluble ammono-basic mercuric nitrate, Hg=N— HgN0 3 which 
is more basic than the foregoing: 

2Hg(N0 3 ) 2 + H 3 .N + 3NH 3 -> Hg 2 (N)N0 3 + 3NH4N0 3 . 

When calomel is treated with ammonium hydroxide, it turns into 
a black, insoluble body. This is a mixture of free mercury, to 
which it owes its dark color, and "infusible white precipitate," 
Hg + Hg(NH 2 )Cl. To this reaction calomel owes its name (Gk. 
KaXofieXns, beautiful black). Mercurous nitrate gives a black, in- 
soluble mixture, 2Hg + Hg2(N)N0 3 . 

Analytical Reactions of Mercury Compounds. — The two 

ionic forms of the element, mercurous-ion Hg + and mercuric-ion 
Hg"^, are both colorless. Their chemical behavior is entirely 
different. Both give the black sulphide HgS, which is insoluble in 
acids and other solvents of mercury salts. Mercurous-ion gives 
the insoluble, white chloride, the black oxide, and a black mixture 
with ammonium hydroxide. Mercuric-ion gives a soluble chloride, 
a yellow, insoluble oxide, and a white precipitate with ammonium 
hydroxide. The behavior with stannous chloride (p. 534) is char- 
acteristic. With potassium iodide the two ions behave differently 
(p. 534). More active metals displace mercury from all com- 
pounds. Copper is used as the displacing metal, in testing for 
Hg + or Hg" 1 " 1 ", because the silvery mercury is easily seen on its 
surface. 

Salts of mercury are volatile. When heated in a tube with 
sodium carbonate, they give a sublimate of metallic mercury. 



RECOGNITION OF CATIONS IN QUALITATIVE ANALYSIS 537 

The Recognition of Cations in Qualitative Analysis 

"Wet-way" analysis consists in recognizing the various positive 
and negative ions present in a solution (p. 436). In discussing 
hydrogen sulphide (p. 273), it was stated that the sulphides might 
be divided into three classes, according to their behavior towards 
water and acids. Now these differences furnish us with a basis for 
distinguishing the cations present in a solution. 

The following plan, taken in conjunction with the statements in 
the context, shows how a single cation may be identified, and how, 
when several cations are present, a separation preparatory to 
identification may be effected. What will be said applies only to 
the case of a solution containing salts like the chlorides, nitrates, 
or sulphates of one or more cations, and leaves the oxalates, 
phosphates, cyanides, and some other salts, out of consideration. 

Group 1. — Add, first, hydrochloric acid, to find out whether 
cations giving insoluble chlorides are present. Argentic, mer- 
curous, and plumbic salts give the white AgCl, HgCl, and PbCl 2 , 
respectively (cf. p. 164). Filtration eliminates the precipitate, if 
there is any. 

Group 2. — A free, active acid being now present, hydrogen 
sulphide is led into the solution. The sulphides insoluble in active 
acids, namely, HgS, CuS, PbS, Bi 2 S 3 , CdS, As 2 S 3 , Sb 2 S 3 , SnS, SnS 2 , are 
therefore thrown down. The first four are black or brown, the next 
two and the last are yellow, and the remaining two are orange 
and brown respectively. A dark-colored substance will naturally 
obscure one of lighter color, if more than one is present. Filtration 
again eliminates the precipitate. 

This group is easily subdivided. Any or all of the last four sul- 
phides will pass into solution when warmed with yellow ammonium 
sulphide, for they give soluble complex sulphides (q.v.). The first 
five sulphides, or any of them, will be unaffected. On the other 
hand, these five sulphides, with the exception of HgS, will interact 
with hot nitric acid (p. 531). Other reactions are then used to 
distinguish between, or, if there is a mixture, to separate, the mem- 
bers of the sub-groups. 

Group 3. — The solution (filtrate) is now neutralized with am- 
monium hydroxide, and ammonium sulphide is added. Some 
ammonium chloride is also used, to prevent the precipitation of 



538 COLLEGE CHEMISTRY 

magnesium hydroxide (p. 525), which, in any event, would be in- 
complete. The sulphides which are insoluble in water, and are 
not hydrolyzed by it, now appear. They are FeS, CoS, NiS, all 
black, MnS and ZnS, which are pink and white respectively. 
There are precipitated also the hydroxides of chromium and of 
aluminium, Cr(OH) 3 and Al(OH) 3 , because their sulphides are 
hydrolyzed by water. 

Group 4. — After filtration, ammonium carbonate is added, and 
precipitates the remaining metals whose carbonates are insoluble, 
BaC0 3 , SrC0 3 , CaC0 3 , with the exception of magnesium (p. 526). 

By addition of ammonium phosphate to a portion of the filtrate, 
magnesium, if present, now comes out in the form NHtMgPO^ 
There remain in solution only salts of potassium, sodium, and 
ammonium. Since only ammonium compounds and other sub- 
stances which can be volatilized have been added, evaporation 
and ignition of the residue leaves the salts of the two metals. 
Salts of ammonium must be sought in a fresh sample by the usual 
test (p. 345). 

Exercises. — 1. Why should we expect ammonium sulphide 
solution to precipitate magnesium hydroxide, and why does it not 
do so? 

2. What volume of air is required to oxidize one formula-weight 
of zinc sulphide to ZnO and S0 2 , and what volume of sulphur 
dioxide is produced? Is the gaseous product more or less diluted 
with nitrogen than when pure sulphur is burned, and by how 
much? 

3. Make equations showing, (a) the effect of heating zinc 
chloride with cobalt nitrate Co(N0 3 ) 2 in the Bunsen flame (p. 
529), (6) the action of hydrogen sulphide on sodium zincate, (c) 
the actions of concentrated nitric acid and of concentrated sul- 
phuric acid on mercury. 

4. What kind of salts might take the place of sodium acetate in 
the precipitation of zinc sulphide (p. 530)? Give examples. 

5. Why do none of the salts of the elements in this family give 
recognizable effects with the borax bead? 



CHAPTER XXXIX 

ELECTROMOTIVE CHEMISTRY 

We have seen that many chemical changes are accompanied by 
a liberation of energy. If no special arrangement is made, the 
energy is always liberated in the form of heat, light, and mechan- 
ical energy. In changes involving ionogens, however, the energy 
can be secured in the form of electricity. Since the change sets 
an electric current in motion, the subject is called electromotive 
chemistry. A knowledge of this branch of the science is essential 
for understanding the numerous commercial applications of 
electricity in chemistry. It also furnishes us with a simple method 
for measuring chemical affinity in ionic reactions. 

Units of Electrical Energy. — Two different units are required 
for defining a quantity of electrical energy. One of these is the 
quantity of electricity, which is expressed in coulombs (p. 237). 
The other is the electromotive force (E.M.F.) of a current, or the 
difference in potential, if a current is not flowing, or the flow is not 
being considered. This is measured in volts. It will be recalled 
that in electrolysis equal quantities of electricity liberate equiva- 
lent weights of the component ions (Faraday's law, p. 231). We 
shall see, however, that with different substances, different differ- 
ences in potential (voltages) are required to produce the de- 
composition. A quantity of electrical energy, used or produced, 
is expressed by the product of the two factors: 

No. of coulombs X No. of volts = Quant, of elect, energy (in Joules). 

If we consider the time occupied by the process, the rate at which 
the electricity flows is expressed in amperes. One coulomb per 
second is one ampere. Hence: 

No. of amperes X No. of volts = Joules per sec. = Watts. 

The kilowatt is 1000 watts. The horsepower is 746 watts. 

An illustration will show the meaning of this relation. If a 
50- watt (16-candle power) incandescent lamp is used on a 110- 

539 



540 COLLEGE CHEMISTRY 

volt circuit, by substituting these values in the equation we per- 
ceive that such a lamp must carry about 0.5 amperes, or one 
coulomb every two seconds. If, with the same voltage, we wanted 
a lamp to carry more electricity per second, we should have to 
reduce the resistance of the lamp, say, by shortening the filament, 
or using a thicker one. Evidently, the number of such lamps re- 
quired to consume one horsepower would be 736/50, or between 
14 and 15 lamps. Again, to decompose one molecular weight of 
hydrochloric acid (36.5 g.) 96,540 coulombs (p. 237) are required, 
and an E.M.F. of at least 1.83 volts (see p. 548). The electrical 
energy needed is therefore 96,540 X 1.83 = 176,670 joules. If 
this were to be accomplished by the current from a 110-volt 
direct-current lighting circuit, passing through a 50-watt lamp in 
series with the electrolytic cell, the time required (x seconds) 
would be given by: 50 joules per sec. X x sees. = 176,670 joules, 
where x = 3533 seconds, or about 59 minutes. 

The factors of electrical energy (volts and amperes) are easily 
measured when electricity is produced, and are easily provided 
according to any specification when electricity is to be used. 
Hence, it is much easier to study the relations between chemical 
change and this form of energy than between the same change 
and the heat or any other form of energy which, under other con- 
ditions, it might produce. Electrochemistry is, therefore, in 
many ways better understood, and easier to handle than are 
other branches of chemistry involving energy. 

Some Reactions that can be Used to Furnish Electricity. 

— A few illustrations of the kinds of reactions which can easily be 
carried out in cells, so as to furnish an electric current instead of 
heat, may be classified thus: 

Combination cells, such as one in which zinc (or some other 
active metal) and bromine are the reacting substances. If zinc 
be placed in bromine-water (or with pure bromine), we obtain 
zinc bromide: 

Zn + Br 2 -> ZnBr 2 , or Zn° + 2Br° -> Zn++ + 2Br~. 

Displacement cells, such as one with cupric sulphate solution 
and a metal more active than copper (e.g., Mg, Al, Zn, or Fe), 
and able to displace (p. 260) this element: 

Zn + CuS0 4 -^ZnS0 4 + Cu, or Zn° + Cu++ -> Zn++ + Cu°. 



ELECTROMOTIVE CHEMISTRY 



541 



A non-metal may also be displaced: 

2KI+Br 2 -+2KBr + I 2 , or I" + Br° -> 1° + Br~. 

Oxidation cell, such as one in which ferrous chloride FeCl 2 or 
stannous chloride SnCl 2 is oxidized by chlorine-water, giving 
FeCl* or SnCU: 

SnCl 2 + Cl a -> SnCl4, or Sn++ + 2C1° -> Sn++++ + 2CT. 

Concentration cells, or cells in which the same substance in two 
different concentrations is used. 



The Arrangement of the Cell. — Every cell has one striking 
characteristic. If the pairs of substances mentioned in the last 
section are placed together, they interact and heat is produced. 
There is no way to avoid the action, and the liberation of the 
energy as heat, if the substances come in contact. If, therefore, 
all the energy is to be obtained as electrical energy, the substances 
must be prevented from com- 
ing in contact with one another. ^^\Ls 
Paradoxical as it may seem, Zn ^^ "\^ Pt 
it is easily possible to obtain 
the electricity, and yet fulfill 
this essential condition. The 
plan in all cells is to place the 
one substance in or round one 
pole, and the other substance 
in or round the other pole, 
and to separate the substances 
by a porous partition, or some 
equivalent arrangement. 

Suppose that it is the first 
of the above-mentioned ac- 
tions that is to be used — the 
action of zinc and bromine. 
The active substances are ar- 
ranged as follows: The pole 
on the left (Fig. 122) is metallic zinc. The solution on the right 
contains the bromine. The porous partition in the center is per- 
meable by migrating ions, but hinders the mere diffusion of the 













Bro 

4-CT+Na+-+ 

Tl 

NaCl 






<-Cr+Na+-> 

Tl 

NaCl 



Pos. ions -*• Neg. ions+- 

Fig. 122. 



542 COLLEGE CHEMISTRY 

dissolved bromine towards the zinc, and so prevents direct inter- 
action with liberation of heat. 

Now, to enable the cell to operate, inactive, conducting sub- 
stances must be added to complete the arrangement. A pole is 
added on the right, a conducting solution is placed to the left of 
the partition, and a wire must connect the two poles. The wire 
may connect the poles through a voltmeter, so that the E.M.F. 
produced may be measured. Also, since bromine-water is a poor 
conductor, a well-ionized salt must be present along with the 
bromine. The substances used for these purposes must be in- 
active. For example, the pole on the right must be a conductor, 
but its material must not interact chemically with the bromine 
or with the salt. A rod of carbon or a platinum wire will serve 
the purpose. A more active metal, such as copper, could not be 
used, because it would combine with the bromine. Again, com- 
mon salt or sodium nitrate may be mixed with the bromine, be- 
cause it will not interact with bromine or carbon or platinum. 
Still again, the solution added on the left must be one which will 
not act upon the zinc pole, or upon the solution on the right, 
which it meets inside the porous partition. Common salt or 
niter fulfills these conditions. An acid could be used on the right, 
but not on the left, for it would interact with the zinc. The 
reader should make a different selection of inactive materials, so 
as to become familiar with the reasoning involved in the choice 
in each case. 

Note that in each figure, the symbols for the active substances 
are in black-face type, the products are in Roman type, and the 
inactive materials are in italic type. 

The Operation of the Cell. — When the cell has been as- 
sembled, and the wires have been connected, the following phe- 
nomena are observed: 

1. The zinc begins to form zinc ions, Zn — > Zn 44 ", an operation 
which leaves the pole negative (Fig. 123). 

2. The bromine molecules nearest their pole touch this pole, 
become bromide ions, Br 2 — » 2Br~, and leave on the pole a positive 
charge. 

3. Since one pole is negative and the other positive, a current 
flows through the wire. 



ELECTROMOTIVE CHEMISTRY 



543 



5 


^y^ 


! 

P 

;. 


^v^Bj + 


Bro x 

Tl 
NaCl 






„ ++ 
— >Zn— > 

<- Cl-+Na+^> 

u 

NaCl 



4. The new positive ions (Zn ++ ) round the left pole (anode) 
attract all the negative ions in the cell, and cause them to migrate 
towards the left so as to keep all parts of the solution neutral. 

5. The new negative ions 
on the right (Br~) similarly 
attract all the positive ions in ^n 
the cell, and cause them to 
drift slowly towards the right 
pole (cathode). 

6 (Very important) . It will 
be seen that the zinc and 
the bromine become ionized 
at a distance from one another 
and do not actually combine. 
The slow migration of the 
Zn" 1-1 " and Br~ ions will, of 
course, after some hours or 
days, bring some of these ions 
together in or near the parti- 
tion, and some molecules will 
be formed. But this operation 
produces no electrical energy 
— it only gives out or absorbs heat (p. 255). It is not an essential 
part of the operation of the cell. The chemical change which pro- 
duces the current is the ionization of the two elements, separately. 
The term combination cell is, therefore, misleading. The cell, 
as a source of electrical energy, is concerned only with producing 
two kinds of ions from the elements. True, these ions, if they 
united, would give the product shown in the equation (ZnBr 2 ), 
but the union, if it ever occurred, would be without electrical 
effect. It is clear that, since there is sodium chloride (or some 
other ionogen) in all parts of the cell, molecules are ionizing, and 
ions are combining, continually, throughout the whole system. 
Thus, on the left some zinc chloride molecules are formed and on 
the right some sodium bromide molecules, and eventually near the 
center some zinc bromide molecules. But these reactions occur 
in every solution containing ionogens, without giving any current. 
In a cell, the only reactions which contribute materially to the 
current are those taking place at the surfaces of the poles. 



Pos. 



ions -*■ 

Fig. 



Neg. ions-*- 



123. 



544 



COLLEGE CHEMISTRY 



A Displacement Cell. — In a similar way, a cell using metallic 
zinc and cupric sulphate solution may be arranged (Fig. 124). 
The zinc forms one pole, and the cupric sulphate solution must be 
placed on the other side of the partition. For inactive materials, 
a plate of copper or of some metal below copper in the activity 
series may be used, and any solution (such as zinc chloride solu- 
tion) which will interact neither with the zinc nor with the cupric 
sulphate. 

In following the operation of the cell, we may start at either 
pole. Thus, the zinc gives zinc-ion Zn° — > Zn ++ -|- 2G. The wire 



Cu Pt 



. 


K^ 


I 

1 


\s 


'+ 


U 

C\xSO t 

_ 






,-*Zn±t 

<—2Ct~+Zn — > 

ti 
ZnCh 






Pos. ions 



Neg. ions*+- 



Fig. 124. 



Pos. ions-*- Neg. ions+- 

Fig. 125. 



becomes negatively charged. The cupric-ion is discharged on the 
other pole CU++ — > Cu° + 2©, rendering it positive. All the 
positive ions in the cell migrate towards the right pole (cathode). 
All the negative ions migrate towards the left pole (anode), since 
positive ions are being formed on the left and are disappearing on 
the right. 

When bromine displaces iodine, the cell may be arranged as 
in Fig. 125. The iodine liberated dissolves in the potassium iodide 
solution and, with starch emulsion present, its formation can be 
detected in a few seconds. 



ELECTROMOTIVE CHEMISTRY 



545 



The Oxidation Cell. — The arrangement whereby stannous- 
ion Sn" 1-1- is oxidized by chlorine-water to stannic-ion Sn 1 ' ' ' is 
shown in Fig. 126. The chlorine Cl° encountering the pole be- 
comes negatively charged, leaving the pole positive. This posi- 
tive charge is shared by the 
whole conducting wire and, 
at the other pole, furnishes 
the positive electricity re- 
quired to raise the charge of 
each tin ion from Sn^ to 
g n ++-H- # Only the tin ions 
which touch the pole can ac- 
quire the charge. 



Facts Concerning all 
Cells. — If the wire is discon- 
nected, the progress of the 
chemical action is stopped, 
although the difference in po- 
tential remains. The charge 
conferred upon a pole, such 
as that from the cupric ion 



b 


i^^ 


7 

J 

.4 

$ 


%3 


j + 


<— ci- r 

<-Cl-+ATa+— > 

IT 
NaCl 






/iSn++++ 

\ 

Sn++ 

<r-2Cl-+Sn++ 

IT 

SnC7 2 



Pos. Ions 



Ney. ions+- 



Fig. 126. 

(Figri24), must be conducted away, before additional charges 
will be transferred to it. 

If a glass partition is substituted for a porous one, the cell 
ceases to generate electricity. The partition must permit the 
trans-migration of the ions, which is a necessary part of the oper- 
ation of the cell. 

When the circuit is closed, the changes described go on until 
one of the active materials is exhausted — for example, until all 
the cupric-ion has been deposited as copper, or until all the zinc 
has been consumed. 

The quantity of electricity producted is 96,540 coulombs for each 
equivalent weight of the active materials transformed, e.g., for 
every 65.4/2 g. of zinc consumed. The rate at which the elec- 
tricity is produced is, in general, greater the larger the area of 
the poles. The amperage of a single cell is, in general, very low. 

The E.M.F. of the cell is not changed by altering the size or 
shape of the poles, or by using more or less of the solutions. It 



546 COLLEGE CHEMISTRY 

is affected by any change in the qualities of the active materials, 
however. Changing the concentration (see p. 551) of the cupric- 
ion (Fig. 124) or of the bromine-water (Fig. 125) has an immediate 
effect. So has substituting one active metal for another (see p. 
547), as magnesium for zinc (Fig. 123). Even hammering the 
metal, thus making it denser, has a slight effect. 

Single Potential Differences Produced by the Metals. — 

If we reconsider the cells described, we shall see that there are 
really two chemical actions in each cell and that these are to some 
extent independent. We can leave the zinc (Fig. 123) constant, 
and change the concentration of the bromine or even substitute 
chlorine or iodine for the latter. On the other hand, we can 
leave the bromine-water constant, and exchange the zinc for some 
other active metal. Thus, the E.M.F. of every cell is really the 
resultant of two effects. Now, these effects can be measured, 
separately. 

If we place zinc in a solution of zinc chloride, we find that then 
is at once a difference in potential between the metal and the solu- 
tion! The metal has an individual tendency to become ionic 
a sort of solution pressure — and to form a few ions, thus making 
the liquid positive and the metal negative. In reality, it is th( 
tendency of the atoms of the metal to give up electrons (p. 235), 
e.g., Zn — 2e = Zn ++ , which is being observed. On the other 
hand, the ions have a tendency to deposit themselves, and a few 
may be deposited (taking up their electrons and becoming neutral), 
rendering the pole positive and the solution negative. If the 
former tendency (the tendency to give up electrons) is the stronger 
of the two (the more active metals), then a difference in potential 
is produced, with the solution positive. If the latter tendency is 
the stronger (less active metals) the solution is observed to be 
negative. Since raising the concentration of the metal-ions will 
increase the tendency to deposition and vice versa, it is customary 
to take as the standard solution, for this purpose, one in which 
the concentration of the metal-ions is normal (N). In the follow- 
ing table, the sign preceding the number is the charge of the 
solution. 



ELECTROMOTIVE CHEMISTRY 



547 



Potential of N Solutions in Contact with Metals 







(E.-M. Series) 






K 


(+ 2.6) 


Fe(Fe++) +0.2 


Bi 


- 0.63? 


Na 


(+ 2.4) 


Cd +0.16 


Sb 


- 0.71? 


Ba 


(+ 2.6) 


Co + 0.05 


Hg(Hg+) 


-0.99 


Sr 


(+]2.5) 


Ni - 0.02? 


Pd 


- 1.03? 


Ca 


(+ 2.4) 


Pb - 0.12 


Ag 


-1.04 


Mg 


+ 1.3 


Sn(Sn++) -0.14 


Pt 


- 1.10? 


Al 


+ 1.0 


H - 0.24 


Au 


-1.7? 


Mn 


+ 0.8 


As - 0.53 






Zn 


+ 0.5 


Cu (Cu++) - 0.58 







Thus, opposite Mg we find +1.3. This means that when a 
piece of magnesium is placed in a solution of a salt of magnesium, 
containing normal concentration of Mg**, the solution is posi- 
tively charged (the metal negatively) and the difference in poten- 
tial is 1.3 volts. With silver in a solution of a salt of silver, 
containing normal concentration of silver-ion, the solution is 
negative and the difference in potential is —1.04 volts. 

For a hydrogen pole, a piece of palladium saturated with hydro- 
gen (p. 57) is used. The values for the metals which decompose 
water with ease cannot be observed, and so calculated values are 
given in parentheses. 

Applications: E.M.F. of a Displacement Cell. — For a cell 
in which one metal is going into solution and another is being 
deposited — like that with zinc and 
cupric sulphate — we can calculate from 
the foregoing data the E.M.F. of the 
cell. Zinc in normal zinc-ion solution 
makes the solution +0.5. Copper in 
normal cupric-ion solution makes the 
solution —0.58. The anodic and cathod- 
ic systems are here stated as if they 
worked against one another. The com — 
bined effect of the two is therefore 
the difference: + 0.5 - (-0.58) = 1.08 
volts. 

The Daniell or gravity cell (Fig. 127) represents this combi- 
nation. The copper plate is at the bottom and the zinc is sus- 





Fig. 127. 



548 COLLEGE CHEMISTRY 

pended above it. The cell is filled with dilute sodium chloride 
solution and crystals of cupric sulphate are thrown in. So long 
as the cell is not disturbed, the heavy, saturated solution of 
cupric sulphate remains at the bottom, so that no porous par- 
tition is required. The actual E.M.F. of this cell is not exactly 
that calculated for normal solutions, because the cupric sulphate 
is in saturated solution, and the concentration of the zinc-ion 
varies, starting at zero and increasing as the cell is used. It is, 
however, a little over 1 volt. 

The Weston Standard Cell contains a pole of mercury in a 
saturated solution of mercurous sulphate and cadmium in contact 
with saturated cadmium sulphate solution. For normal solu- 
tions, the voltage would be +0.16 - (-0.99) = 1.15 volts. At 
20° it is 1.0183 volts. 

The Clark Standard Cell contains zinc and zinc sulphate solution 
in place of the cadmium. With normal solutions it would give 
+ 0.5 - (-0.99) = 1.49 volts. It actually gives 1.434 volts. 

Single Potential Differences for Non-Metallic Ions, — 

The corresponding figures for the non-metals are: 

I -0.78 Br -1.32 CI -1.59 

Hence the cell with zinc and bromine-water (p. 541) in presence 
of normal concentration of the respective ions gives +0.5 — 
(—1.32) = 1.82 volts. Similarly, the cell in which bromine dis- 
places iodine (p. 544) gives -0.78 - ( — 1.32) = 0.54 volts. 

Applications: Electrolysis: Discharging Potentials. — 

When a solution of a salt, such as cupric chloride, is electrolyzed, 
copper and chlorine are liberated at the two poles. Now, when 
the electrolysis has made some progress, if the battery is taken 
out, and the wires are joined, a current, the polarization current, 
flows. Evidently, the copper and chlorine liberated in and round 
the electrodes have made the arrangement into a copper-chlorine 
battery cell. Assuming normal concentrations, the E.M.F. of 
the polarization current is —0.58 — ( — 1.59) = 2.17 volts. Now 
this counter-current is in operation during the whole electrolysis. 
To overcome it, and maintain the electrolysis, evidently an 
E.M.F. of at least 2.17 volts from the battery is required. This 
is called the discharging potential for cupric chloride. 



ELECTROMOTIVE CHEMISTRY 549 

Applications: Electrolytic Refining, — The electrolytic 
process for refining copper (read p. 511) can now be more easily 
understood. Both electrodes are made of copper, and the solu- 
tion contains cupric sulphate. There is, therefore, no difference 
in potential between the plates, except a very small one, due to 
the fact that one plate is pure copper and the other impure. 
Hence a very slight E.M.F., sufficient to overcome the difference 
just mentioned, and to overcome the friction of the moving ions, 
is all that is required, and 0.5 volts is sufficient. 

As regards the resulting purification, the anode of crude copper, 
which is being consumed, contains, besides copper, small amounts 
of less active metals like silver and gold, and of more active 
metals like zinc. So far as the more active metals are concerned, 
the cell is like one with zinc and cupric sulphate (p. 544). It 
would run by itself, without any outside current, and would 
actually generate a current. Hence the active metals become 
ionic easily, and displace cupric-ion from the solution. The less 
active metals, on the other hand, are not required for the trans- 
ference of the electricity, since a great excess of the more active 
copper is available. They also require a larger E.M.F. for their 
ionization than does copper. Hence, they remain as metals, and 
drop to the bottom of the cell (sludge) as the anode of crude 
copper wears away. 

Applications: Couples. — The fact that metallic zinc will 
displace hydrogen-ion from an acid, or cupric-ion from cupric 
sulphate solution can now be explained. The more active metals 
are the ones which have the greatest tendency to become ionic. 
Each will deprive the ions of a metal below it in the list of their 
electric charges: 

Zn° + 2H+->Zn++ + H 2 °T. 
Zn o + Cu++ -» Zn++ + Cu°| . 

Now we have noted the facts (pp. 55, 626) that contact with a 
platinum wire, or the presence of impurities (other metals) in 
the zinc, will hasten its action. Pieces of two metals in contact 
with one another constitute a couple. With zinc and platinum 
in an acid, a current is set up, like that of a short circuited cell. 
The zinc becomes negative, the platinum positive, and the hydro- 
gen is liberated upon the platinum. This facilitates the action 



550 COLLEGE CHEMISTRY 

because, when the platinum is absent, and the hydrogen gas, in bub- 
bles, is liberated on the surface of the zinc, this surface is only 
partly in contact with the acid (H+), and so the liberation of the 
hydrogen is slower. 

Galvanized iron is also a couple. When rain (dilute carbonic 
acid) falls upon it, the zinc, being the more active metal (p. 528), 
is the anode and tends to become ionized (forming the carbonate) . 
The iron is the cathode and is not affected. The carbonate, how- 
ever, forms a closely adhering coating on the zinc, and so but 
little of this metal is actually consumed, and the material is 
therefore durable. On the other hand, a sheet of iron, without 
the zinc coating, gives ferrous carbonate which is easily oxidized 
to ferric hydroxide (a base too weak to give a carbonate). This 
forms a brittle, porous layer which does not mechanically protect 
the surface from further action, and so the iron is finally all oxi- 
dized. Tin-plate (tin on iron, a couple) is not attacked so long 
as the layer of tin is nowhere broken. But damaged tin-plate 
rusts rapidly. There, the iron is the more active metal (p. 547) 
and forms carbonate and then hydroxide continuously, while 
the tin remains unaffected. 

Applications: Measurement of Affinity. — Since equal 
quantities of electricity bring about (or are brought about by) 
chemical changes in chemically equivalent weights of material, 
it follows that the E.-M. forces required (or produced) are pro- 
portional to the chemical affinity. Thus the activities of the 
metals, expressed in volts (p. 547), are accurate figures for the 
relative affinities of the metals, so far at least as ionic actions are 
concerned. In point of fact, they express also the approximate 
affinities of the metals in other actions (pp. 60, 531) as well. Again, 
by using different oxidizing agents in place of the chlorine-water 
(p. 545) and noting the differences in potential, we can obtain 
numbers representing the relative activities of various oxidizing 
agents towards oxidizable ions. 

Concentration Cells. — If two rods of a metal (e.g., tin) are 
placed together in the same solution of a salt of the metal (e.g., 
stannous chloride SnCl 2 ), there is no difference in potential, be- 
cause the state of both poles is in all respects the same. But, if 



ELECTROMOTIVE CHEMISTRY 



551 



the solution round one pole is more concentrated than that 
round the other, a difference in potential is produced (Fig. 128). 
The tendencies of the metallic tin to form ions are equal, but 
the pressures of the stannous 
ions are different, and so, when 
the circuit is closed, stannous Sn 
ions are discharged on the tin + | - 
pole in the more concentrated 
solution, forming long crystals 
of tin, and tin in equal amount 
from the pole in the dilute solu- 
tion becomes ionic. 

The law which formulates 
the relation between the two 
concentrations and the E.M.F. 
produced being known, it is 
possible to use the concentra- 
tion cell for measuring solu- 
bilities of insoluble salts. 
Thus, we cannot easily meas- 
ure the solubility of silver chlo- 
ride by the ordinary method 
(p. 123), because evaporation of the solution may leave a larger 
mass of impurities, derived from solution of the glass, than of dis- 
solved silver chloride. Hence, we use two poles of silver, place one 
in normal silver nitrate solution and the other in saturated silver 
chloride solution (with excess of the solid), measure the difference 
in potential, and calculate the ratio of the concentrations of silver- 
ion in the two solutions. The absolute value of that in the silver 
nitrate solution is known, and so the absolute value of the Ag + 
concentration in the silver chloride solution can be found. Since 
silver chloride is a salt (p. 242), it is very highly ionized in so 
dilute a solution, and the molecular concentration of silver-ion is 
practically equal to the total molecular concentration of silver, 
and therefore of silver chloride in the liquid. 




Pos. ions-* Meg- ions-*- 



Fig. 128. 



Exercises. — 1. Make diagrams of the following cells, choosing 
with care suitable inactive substances to complete the arrange- 
ment: (a) chlorine-water and aluminium; (6) chlorine- water 



552 COLLEGE CHEMISTRY 

and ferrous chloride; (c) zinc and dilute sulphuric acid; (e) chlo- 
rine-water and potassium bromide. 

2. Calculate the E.M.F. of each of the cells in Ex. 1, assuming 
normal solutions to be present. 

3. What will be the discharging potentials of solutions of the 
following substances, assuming N concentrations of the ions: 
(a) manganous chloride; (6) hydrogen iodide; (c) ferrous bromide; 
(e) sodium chloride (hydrogen is liberated)? 

4. What weight of zinc must be ionized every hour in a cell in 
order to produce a current of 5 amperes strength? For how long 
would 500 g. of zinc serve to maintain this current? 

5. In the zinc-bromine cell (p. 543), why is the zinc pole called 
the anode, although its charge with respect to the platinum is 
negative? 



CHAPTER XL 

ALUMINIUM AND THE METALS OF THE EARTHS 

The chief members of the family occupying the fourth column of 
the periodic table are: boron (B, at. wt. 11), aluminium (Al, at. wt. 
27.1), gallium (Ga, at. wt. 70), indium (In, at. wt. 115), thallium 
Tl, at. wt. 204), all on the right side of the column; and scandium 
(Sc, at. wt. 44.1), yttrium (Y, at. wt. 89), lanthanum (La, at. wt. 
139), on the left side. These elements are all trivalent. 

The Rare Elements of this Family. — The oxide and hydrox- 
ide of boron are acidic (p. 431). Those of aluminium (Al(OH) 3 ), 
gallium (Ga(OH) 3 ), indium (In(OH) 3 ), and thallium (TIO.OH) are 
basic, but behave also as acids towards strong bases. 

Gallium and indium occur occasionally in zinc-blende, and were 
discovered by the use of the spectroscope. The former takes its 
name from the country (France) in which the discovery was made, 
and the latter from two blue lines shown by its spectrum. 

Thallium is found in some specimens of pyrite and blende. It 
was discovered by Crookes, by means of the spectroscope, in the 
seleniferous deposit from the flues of a sulphuric acid factory. It 
received its name from the prominent green line in its spectrum 
(Gk. OaXXos, a green twig). It gives two complete series of com- 
pounds. In those in which it is trivalent (thallic salts), it resem- 
bles aluminium (q.v.) . Thus, the salts of this series are more or less 
hydrolyzed by water. Univalent thallium recalls both sodium and 
silver. Thallous hydroxide TIOH is soluble, and gives a strongly 
alkaline solution, but the chloride is insoluble in cold water. The 
solutions of the thallous salts are neutral. The metal is displaced 
from its salts by zinc. 

Of the elements on the left side of the column, scandium, whose 
existence and properties were predicted by Mendelejeff (p. 301), is 
the best known. The metals of the rare earths, of which it is one, 
are found in rare minerals such as euxenite, gadolinite, orthite, and 

553 



554 COLLEGE CHEMISTRY 

monazite, which occur in Sweden, Greenland, and the United 
States. Cerium (Ce, at. wt. 140.25), neodymium (Nd, at. wt. 
144.3), and praseodymium (Pr, at. wt. 140.9) occur along with 
lanthanum in cerite, a silicate of these four elements. These four 
are included amongst the metals of the rare earths. The com- 
pounds of many of these rare elements behave so much alike that 
separation is difficult. There are several with atomic weights 
near to that of lanthanum for which accommodation cannot easily 
be found in the periodic table. Ostwald has compared them to a 
group of minor planets such as in the solar system takes the place 
of one large planet. 

Aluminium 

The Chemical Relations of the Element. — Aluminium is 
trivalent exclusively. Its hydroxide, like that of zinc (p. 529), is 
amphoteric, that is to say, it is feebly acidic as well as basic, and 
hence the metal forms two sets of compounds of the types Na3.A103 
(sodium aluminate) and A1 2 .(SC>4) 3 . The salts of both series are 
more or less hydrolyzed by water, the former very conspicuously 
so. It is worth noting that the hydroxides of the trivalent metals, 
or metals in the trivalent condition, such as Al(OH) 3 , Cr(OH) 3 , 
Fe(OH) 3 , are all distinctly less basic than are those of the bivalent 
metals such as Zn(OH) 2 , Cd(OH) 2 , Fe(OH) 2 , Mn(OH) 2 . Alumin- 
ium does not enter into complex anions or cations, and is too feebly 
base-forming to give salts like the carbonate or sulphide. 

Occurrence. — Aluminium is found very plentifully in combi- 
nation, coming next to oxygen and silicon in this respect. The 
felspars (such as KAlSi 3 8 ), the micas (such as KAlSi0 4 ), and 
kaolin (clay) H 2 Al 2 (Si0 4 ) 2 ,H 2 are the commonest minerals con- 
taining it. Since the soil has been formed largely by the weather- 
ing of minerals like the felspars, clay and other products of the 
decomposition of such minerals constitute a large part of it. Cryo- 
lite is a double fluoride 3NaF,AlF 3 . Various forms of the oxide 
and hydroxide are also found. 

Preparation and Physical Properties. — The metal is now 
made on a large scale by electrolysis of the oxide A1 2 3 dissolved 
in a bath of molten cryolite (m.-p. 1000°), a process invented by 
C. M. Hall (1886). The operation is conducted in cells (5x3 feet, 




ALUMINIUM 555 

or larger), the carbon linings of which form the cathodes (Fig. 129). 
The anodes are rods of carbon which combine with the oxygen as 
it is liberated. The molten metal (m.-p. 659°) sinks to the bottom 

of the cell and is drawn off periodically, f 

while fresh portions of the oxide are 
added from time to time. The oxide is 
made from bauxite (see below), and 
must be free from oxide of iron and 
other impurities, as the metal cannot 
be purified commercially. The current 
(E.M.F. 5-6 volts) maintains the tem- 
perature of the molten materials, and 
causes the decomposition. In 1866 FlG - 129 - 

aluminium cost $250-750 (£50-150) per kilogram. In 1883 the 
whole production was about 40 kilos. In 1913 the United States 
alone consumed 35 million kilos, costing about 50 cents (2/-) per kilo. 
The metal melts at 658.5°, but is not mobile enough to make 
castings. It is exceedingly light (sp. gr. 2.6), and in tensile strength 
excels the other metals, with the exception of iron and copper. 
It is malleable, and the foil is taking the place of tin foil for wrap- 
ping foods. It has a silvery luster, and tarnishes very slightly, 
the firmly adhering film of oxide first formed protecting its surface. 
Although, comparing cross-sections, it is not so good a conductor 
of electricity as is copper, yet weight for weight it conducts better. 
It is difficult to work on the lathe or to polish, because it sticks to 
the tools, but the alloy with magnesium (about 2 per cent) called 
magnalium has admirable qualities in these respects. Aluminium 
bronze (5-12 per cent aluminium with copper) is easily fusible, has 
a magnificent golden luster, and possesses mechanical and chemical 
resistance exceeding that of any other bronze. The metal and its 
alloys are used for making cameras, opera-glasses, cooking utensils, 
and other articles requiring lightness and strength, as well as for 
the transmission of electricity. The powdered metal, mixed with 
oil, is used in making a silvery paint. 

Chemical Properties, — The metal displaces hydrogen from 
hydrochloric acid very easily. It displaces hydrogen also from 
boiling solutions of the alkalies, forming aluminates: 

2A1 + 6NaOH -* 2Na 3 A10 3 + 3H 2 . 



556 COLLEGE CHEMISTRY 

In consequence of its very great affinity for oxygen, aluminium 
displaces from their oxides the metals below magnesium in the 
E.-M. series. Thus, when a mixture (thermite) of aluminium 
powder and ferric oxide is placed in a crucible and ignited by means 
of a piece of burning magnesium ribbon, aluminium oxide and iron 
are formed: 

Fe 2 3 + 2A1 -* A1 2 3 + 2Fe. 

The very high temperature (about 3000°) produced by the action is 
sufficient to melt both the iron (m.-p. 1530°) and the oxide of 
aluminium (m.-p. 2050°). The products, not being miscible, 
separate into two layers. This very simple method of making pure 
specimens of metals like chromium, uranium, and manganese, whose 
oxides are otherwise hard to reduce, is called by Goldschmidt, 
the inventor, aluminothermy. By preheating the ends of steel rails 
with a gasoline torch, firing a mass of thermite in a crucible above 
the joint, and allowing the iron to flow into the joint, perfect welds 
are made. In the same way, large castings, like propeller shafts, 
when broken, can be mended. The sulphides, such as pyrite, are 
reduced with like vigor by aluminium. 

The largest part of the aluminium of commerce is used by steel- 
makers. When added in small amount (less than 1 : 1000) to 
molten steel, it combines with the gases, and gives sound ingots 
free from blow holes. 

Aluminium Chloride AlCl 3 . — If the metal or the hydroxide 
is treated with hydrochloric acid, and the solution is allowed to 
evaporate, the hydrated chloride A1C1 3 ,6H 2 is formed. When 
heated, this hydrate is completely hydrolyzed, hydrochloric acid 
is given off, and only the oxide remains. The anhydrous chloride 
AICI3 is made by passing dry chlorine over aluminium. Since it 
sublimes as a white crystalline solid without melting, when thus 
prepared it is vaporized and condenses in a cool part of the tube. 
It fumes when exposed to moist air on account of the hydrogen 
chloride produced by hydrolysis, and only with excess of hydro- 
chloric acid does it give a clear solution free from basic salts. 

Aluminium Hydroxide and the Aluminates. — When an 
alkali is added to a solution of a salt of aluminium, the hydroxide 
Al(OH) 3 is precipitated in gelatinous form. It loses water gradu- 



ALUMINIUM 557 

ally when dried, forming no intermediate hydroxides, until AI2O3 
remains. Natural forms of this substance are hydrargyllite 
Al(OH) 8 (= A1 2 3 ,3H 2 0), bauxite Al 2 0(OH) 4 (= A1 2 3 ,2H 2 0), 
which always contains ferric oxide, and diaspore AIO.OH ( = A1 2 3 ,- 
H 2 0). 

Commercially, the hydroxide is made by heating bauxite with 
dry sodium carbonate, and extracting the sodium metaluminate 
with water: 

Al 2 0(OH) 4 + Na 2 C0 3 -» 2NaA10 2 + C0 2 + 2H 2 0. 

The iron, present as an impurity, remains, as ferric oxide, undis- 
solved. The hydroxide is then precipitated by passing carbon 
dioxide through the solution: 

2NaA10 2 + C0 2 + 3H 2 ->Na 2 C0 3 + 2A1(0H) 3 . 

Aluminium hydroxide, being amphoteric, interacts both with 
acids and with bases, and is, therefore, like zinc hydroxide (p. 531), 
ionized both as a base and as an acid. It interacts only slightly 
with ammonium hydroxide, because this substance is too feebly 
basic, but, from the solution in the active alkalies, the aluminates 
Na 3 .A10 3 , Na.A10 2 , and K.A10 2 , can be obtained in solid form. 
The aluminates are largely hydrolyzed by water: 

NaA10 2 + 2H 2 <=± NaOH + Al(OH) 3 . 

When calcium chloride is added to a solution of sodium alumi- 
nate, the insoluble calcium metaluminate is deposited: 

2NaA10 2 + CaCl 2 -> Ca(A10 2 ) 2 + 2NaCl. 

The relations of these various substances are shown by the follow- 
ing formulae : 

/0 -H /0 -Na Q Q ^ 

Al-O-H Al-O-Na Alf Alf ^Ca 

^O-H ^O-Na °- H °- Na A ,/° 

A number of insoluble metaluminates, such as spinelle Mg(A10 2 ) 2 , 
and gabnite Zn(A10 2 ) 2 , are found in nature. They contain bi- 
valent metals in place of the calcium in the last-named compound. 



558 COLLEGE CHEMISTRY 

Aluminium Oxide Al 2 3 . — The oxide (alumina) is manu- 
factured by heating the pure hydroxide made from bauxite (see 
above). It is found in nature in pure form as corundum. This 
mineral is only one degree less hard than the diamond. Emery 
is a common variety, contaminated with ferric oxide, and was 
widely used as an abrasive, until largely displaced by carborundum. 
The ruby is pure aluminium oxide tinted by a trace of a compound 
of chromium, while the sapphire is the same material colored with 
aluminates of iron and titanium. It is said, however, that the 
same tint is conferred upon colorless corundum by exposure to the 
influence of salts of radium. By ingenious methods of fusing the 
oxide, " synthetic" sapphires and rubies are now made in large 
quantities. Alundum, a refractory material for crucibles, is made 
by heating objects made of the oxide in the electric furnace until 
a small proportion of the material is melted. 

Aluminium Sulphate: The Alums, — Aluminium sulphate 

A1 2 (S0 4 ) 3 ,18H 2 is prepared by treating either bauxite, or the pure 
hydroxide made from bauxite, or pure clay (kaolin) with sulphuric 
acid. In the latter case the insoluble residue of silicic acid is 
removed by filtration: 

H 2 Al 2 (Si0 4 ) 2 + 3H 2 S0 4 -> A1 2 (S0 4 ) 3 + 2H 2 Si0 3 + 2H 2 0. 

The solution of the sulphate is acid in reaction. It crystallizes in 
leaflets which, when the source was clay or bauxite, have a yellow 
tinge due to the presence of iron as an impurity. The salt is used 
as a source of precipitated aluminium hydroxide in paper-making, 
water purification, and dyeing. 

When sulphate of potassium solution is added to a strong solu- 
tion of aluminium sulphate, octahedral crystals of potash alum 
(see below) are deposited. This is a double salt, and is one of a 
large class known as the alums. The alums have the general 
formula ikf 2 I S0 4 ,M 2 III (S0 4 )3,24H 2 0, and may be made as above by 
using a sulphate of a univalent metal with one of a trivalent metal. 
Thus, for M 1 we may use K, NH 4 , Rb, Cs, and Tl 1 , and for 
ilf 111 , Al, Fe m , Cr m , Mn m , and Tl m . All the alums crystallize 
in octahedra. 

Potassium-aluminium sulphate K 2 S0 4 ,A1 2 (S0 4 ) 3 ,24H 2 0, ordinary 
alum, is made from aluminium sulphate. It is also prepared by 



ALUMINIUM 559 

heating alunite, a basic alum found near Rome and in Hungary, 
and extracting the product with water. The alunite KA1 3 (0H)6- 
(S0 4 )2 leaves an insoluble residue of the hydroxide, mixed with 
ferric oxide which is present as an impurity: 

2KA1 3 (0H) 6 (S0 4 )2 -> K 2 S0 4 ,A1 2 (S0 4 )3 + 4A1(0H) 8 . 

The hydrated salt melts at 90°. An aqueous solution of this 
salt, or of sodium phosphate (p. 464), is used for fireproofing 
draperies. The crystals deposited in the fabric melt easily, and 
the fused material protects the fibers from access of oxygen. When 
heated more strongly alum loses its water of hydration, together 
with some sulphur trioxide, and leaves a slightly basic, anhydrous 
salt known as burnt alum. Potash-alum and ammonium-alum are 
more easily freed from impurities (e.g., compounds of iron) by 
recrystallization than is aluminium sulphate, and the alums are 
therefore used instead of the latter in medicine, in dyeing delicate 
shades, and to replace cream of tartar in baking powder (p. 463). 
In the last case, the reaction 

K 2 S0 4 ,A1 2 (S0 4 ) 3 ,24H 2 + 6NaHC0 3 -> K 2 S0 4 + 3Na 2 S0 4 
+ 2A1(0H) 3 + 6C0 2 + 24H 2 

liberates carbon dioxide by hydrolysis of the aluminium carbonate. 

Hydrolysis of Aluminium Carbonate. — The foregoing 

action, and others discussed above (p. 557), show that the carbon- 
ate is completely hydrolyzed: 

A1 2 (C0 3 ) 3 + 6H 2 <± 2A1(0H) 3 j + 3H 2 C0 3 -> 3H 2 + 3C0 2 1 . 

It will be seen that this may be due only in part to the feebly basic 
qualities of the hydroxide. If the hydroxide were not precipitated, 
it would cause some reversal of the action, and some of the car- 
bonate would remain. The insolubility of one product explains 
also other cases of the complete hydrolysis of salts (e.g., ammonium 
silicate p. 429 and next section). 

Aluminium Sulphide A1 2 Sq. — This salt is most easily 
obtained by mixing pyrite with aluminium powder and igniting 
with magnesium ribbon (p. 556) : 

3FeS 2 + 4A1 -> 2A1 2 S 3 + 3Fe. 



560 COLLEGE CHEMISTRY 

It forms a grayish-black solid, and is decomposed by water, as is 
magnesium sulphide, giving the hydroxide and hydrogen sulphide. 
On this account, the sulphide, like magnesium sulphide (p. 538), 
cannot be formed by precipitation in presence of water. Thus, 
ammonium sulphide with a salt of aluminium, in solution, gives a 
precipitate of aluminium hydroxide: 

A1 2 (S0 4 )3 + 3(NH 4 ) 2 S + 6H 2 -> 2A1(0H) 3 [ + 3(NH4) 2 S0 4 + 3H 2 S. 

Coagulation Method of Purifying Water: Sizing Paper. — 

When aluminium hydroxide is formed, it gives first a colloidal sus- 
pension, which coagulates to a gelatinous precipitate. When this 
precipitate is produced in water used for domestic purposes, and 
containing fine, suspended matter, the gelatinous material causes 
the fine particles to collect into larger ones which settle rapidly, 
and permits the use of relatively small settling ponds. These 
larger particles enclose also practically all the bacteria. If the 
water is slightly hard, crude aluminium sulphate, alone, is added: 

3Ca(HC0 3 ) 2 + A1 2 (S0 4 ) 3 -> 3CaS0 4 + 2A1(HC0 3 ) 3 , (1) 
A1(HC0 3 ) 3 + 3H 2 -> Al(OH) 3 J, + 3H 2 C0 3 . (2) 

If the water is very soft, a little lime Ca(OH) 2 is added. The few 
remaining bacteria are destroyed by later addition of bleaching 
powder or chlorine-water (p. 312). 

In some localities crude ferrous sulphate, obtained as a by- 
product in cleaning iron, is cheaper, and is employed instead. The 
lime precipitates ferrous hydroxide Fe(OH) 2 . This is quickly 
oxidized to colloidal ferric hydroxide Fe(OH) 3 , which coagulates 
the suspended matter. 

Aluminium hydroxide is employed also in sizing cheaper grades 
of paper (p. 402), an operation required to prevent the absorption 
and consequent spreading of the ink. For writing-paper, gelatine 
solution is employed. In making printing-papers, rosin soap 
(made by dissolving rosin in caustic soda) is mixed with the paper- 
pulp, and aluminium sulphate is added. The rosin and aluminium 
hydroxide are precipitated in the pulp, perhaps in feeble combina- 
tion, and pressing between hot rollers afterwards melts the former 
and gives a surface to the paper. 

Delicate cloth goods are rendered waterproof by saturating 
them with aluminium acetate solution and then steaming them to 



ALUMINIUM 561 

promote hydrolysis. The aluminium hydroxide is thus precipi- 
tated in the capillaries of the cotton or linen and renders them 
non-absorbent : 

A1(C0 2 CH 3 ) 3 + 3H 2 <=> A1(0H) 8 + 3HC0 2 CH 3 . 

Kaolin and Clay: Earthenware and Porcelain. — By the 

action of water and carbon dioxide upon granite, and other rocks 
containing felspar KAlSi 3 8 , the potash is slowly removed, and 
the compound is changed largely into a hydrated orthosilicate 
H 2 Al2(Si0 4 )2,H 2 0. When pure, it forms kaolin or china clay, a 
white, crumbly material. When washed away and redeposited, it 
usually acquires compounds of iron, the carbonates of calcium and 
magnesium, and sand (silica), becoming common clay. Ocher, 
umber, and sienna are clays colored with oxides of iron and man- 
ganese. Fuller's earth is a purer variety. 

The plasticity of clay, a property connected with the colloidal 
nature of the kaolin, enables it to be fashioned into various shapes. 
When heated, it shrinks and becomes a hard, solid, porous mass, 
and does not melt. These two properties enable us to use it in 
making bricks, pottery, and porcelain. The presence of calcium 
and magnesium compounds makes the clay more fusible, because 
it permits the formation of fusible silicates of these metals. Bricks 
and tiling for roofs and drains are made of common clay and, when 
red, owe their color to oxide of iron Fe 2 3 . The firing is done with 
fuel gas in ovens or kilns of brickwork. To glaze drain pipes and 
some bricks, salt is thrown into the kiln. The water vapor, at a 
red heat, hydrolyzes the salt, hydrogen chloride is set free, and the 
sodium hydroxide gives with the clay a fusible sodium-aluminium 
silicate which fills the surface pores. Clay for fire brick (infusible) 
must contain silica, but no lime. 

China and porcelain are made from pure clay, free from iron, 
to which a little of the more fusible felspar is added. After the 
first firing, the articles are porous (bisque), and must be covered 
with a glaze. A paste of finely ground felspar and silica, some- 
times containing lead oxide, is spread on the surface, and thi 
articles are fired again, at a higher temperature. Colored decora- 
tions are added by means of suitable materials, mainly oxides of 
metals which give colored silicates. The third firing causes these 
oxides to interact and fuse with the glaze. 



562 COLLEGE CHEMISTRY 

The Schwerin process for separating ferric oxide Fe 2 3 from 
clay, so that white porcelain may be obtained, is now used on a 
large scale and affords an interesting application of the properties 
of colloidal suspensions (p. 417). When the impure clay is sus- 
pended in water, the particles of ferric oxide are positively charged 
and those of the clay are negative. By inserting plates connected 
with the dynamo in the trough, the clay particles are caused to 
drift towards the positive plate and the ferric oxide towards the 
other, so that, when the liquid from the positive end is allowed to 
settle, pure clay is obtained. 

In making bricks, in some cases, advantage is taken of the fact 
that negative colloids, such as clay, become more strongly negative 
in presence of a trace of free alkali. Thus, when a trace of sodium 
hydroxide is added to clay slip, the particles repel one another more 
strongly, the cohesion which causes the plasticity is reduced, and 
the clay can be poured into molds. This avoids diluting the clay 
with water, which would only have to be driven out again, with 
great waste of heat, in the firing. 

Cement. — Cement is made by heating limestone CaC03, 
and clay HAlSi04, or a natural rock containing both materials 
in the right proportions. Such a rock, made into cement by 
volcanic heat, was quarried by the Romans near Naples and else- 
where, and its capacity for hardening even under water was utilized 
by them. Blast-furnace slag, when pulverized and heated with 
limestone, has been found to yield an excellent quality of cement, 
and a valuable use has thus been found for what was formerly an 
annoying encumbrance. The mixture, or pulverized natural rock, 
is moistened and fed slowly in at the upper end of an inclined (6°) 
revolving cylinder of iron (20 to 45 by 2 meters). The motion 
continually turns over the thin layer, and exposes every particle 
to the heat of the air-blast, charged with pulverized coal, burning 
in the interior. The product slides out in a continuous stream at 
the lower end, and is pulverized by steel balls in a ball mill. 

Cement is held to be a mixture of calcium silicate and calcium 
aluminate. The former is simply a filler. The latter is hydrolyzed 
by the water: 

Ca 3 (A10 3 ) 2 + 6H 2 -> 3Ca(OH) 2 + 2H 3 A10 3 . 



ALUMINIUM 



563 



The calcium hydroxide slowly crystallizes, connecting the particles 
of the calcium silicate. The aluminium hydroxide fills the inter- 
stices and renders the whole compact and impervious. 

Ultramarine. — Formerly, pulverized lapis lazuli, a rare 
mineral of beautiful blue color, was used by artists as a pigment. 
Gmelin (1828) found a way of making it artificially. A mixture of 
kaolin, sodium carbonate, sulphur, and charcoal is heated until a 
green mass is obtained. This mass is then pulverized and heated 
with more sulphur. The product is used as laundry blueing, in 
making blue-tinted paper, and with oil in making paint. It is also 
added in small amount to correct the natural yellow tint of linen, 
starch, sugar (p. 405), and paper-stock. By varying the mode of 
heating, without altering the composition, various colors from 
green to reddish violet can be obtained. No pure colored sub- 
stance can be extracted from it. The variety of colors is due to 
different degrees of colloidal dispersion of some substance sus- 
pended in the solid, just as gold, which is pale yellow in mass, gives 
colloidal suspensions (p. 416) of different colors (red, purple, or 
blue) according to the fineness of the particles (c/. p. 494). 



Dyeing, — The problem of the dyer is to confer the desired 
color upon a fabric made, usually, of cotton, linen, wool, or silk, 
and to do this in such a way A B 

that the dye is fast to (i.e., is 
not removed or destroyed by) 
rubbing and fight, and often, 
also, to washing with soap. To 
understand the means by which 
this is achieved, it must be noted 
that cotton and linen consist of 
smooth hollow fibers (Fig. 130A) 
of the composition of cellulose 
(C 6 Hio0 5 )x. Wool is made of hollow fibers, with a scaley surface 
(B) and silk of solid filaments, but these are composed of proteins 
(p. 422). Now, the proteins are much more active chemically 
than is cellulose, and also, as colloidal materials, seem to have a 
much greater tendency to adsorb other substances than has cellu- 
lose. Hence, accidental stains on wool or silk are much less often 




U) 



Fia. 130. 



564 COLLEGE CHEMISTRY 

removable than are those on cotton, and when samples of the three 
materials are dipped in a solution of a dye, the first two are per- 
manently dyed, while from the last most dyes can be completely 
washed out with water. 

Three modes of dyeing may be mentioned: 

1. Insoluble Dyes. If the colored body can be produced by 
precipitation, after the solution has filled the capillary and wall 
of every fiber of the goods, then, if the dye is sufficiently insoluble, 
it is mechanically imprisoned in every fiber and cannot be washed 
out. This plan may be applied to any kind of goods. For 
example, if cotton, silk, or wool is first boiled in a solution of lead 
acetate, and is then soaked in a boiling solution of potassium 
chromate, it is dyed a brilliant, permanent yellow. Lead chro- 
mate is the colored body: 

Pb(C0 2 CH 3 ) 2 + K 2 Cr0 4 ±+ 2K(C0 2 CH 3 ) + PbCr0 4 | . 

The part precipitated on the outside of the goods can be, and is, 
at once washed off by rubbing in water, but the particles inside the 
fibers can come out only by being dissolved, and they are insoluble 
in water. Indigo Ci 6 Hi N 2 O 2 , which is used in larger amounts than 
any other dye, belongs to this class. Obtained in early times from 
several plants in Europe and Egypt, where it was known as woad, 
and more recently imported from India, where the cultivation of 
the indigo plant was as important an industry as is the growing 
of cotton in the Southern States, it is now almost all made 
artificially. Synthetic indigo is manufactured, with naphthalene 
CioH 8 (p. 411), obtained from gas tar and the tar from by-product 
coke ovens, as the initial substance. The cloth is saturated with 
an alkaline solution of indigo white Ci6Hi2N 2 2 , a soluble, slightly 
acid substance, and the oxygen of the air subsequently oxidizes 
this and deposits the insoluble indigo blue within the fibers: 
201611^202 + 2 -> 2Ci 6 H 10 N 2 O 2 1 + 2H 2 0. 

Indanthrene blue is applied in the same way as indigo, and is even 
less affected by light. 

2. Mordant or Adjective Dyes. Since cotton is inactive chemi- 
cally and, although a colloid, has but a slight tendency to adsorb 
dyes, it is usually, necessary first to introduce into the fibers of 
cotton some colloidal substance with greater adsorptive powers. 
Substances of this kind are tannic acid for basic dyes, and gelati- 



ALUMINIUM 565 

nous colloidal hydroxides, such as those of aluminium, tin, iron and 
chromium, for non-basic (including acid) dyes. They are called 
mordants (Lat. mordere, to bite). Thus, if in three jars we place 
very dilute solutions of aluminium sulphate, ferric chloride FeCla 
and chromous acetate Cr(C0 2 CH 3 ) 2 , then add a few drops of a 
solution of a dye to each, and finally introduce a little of a base 
(like sodium hydroxide) to precipitate the hydroxide of the metal, 
this hydroxide will adsorb the dye and carry it into the precipitate. 
Such a precipitate of mordant and dye is called a lake. With the 
same dye, the three lakes have different colors. Thus, in the above- 
mentioned experiment, if alizarin (madder) is used as the dye, the 
colors are red (Turkey red), violet, and maroon, respectively. 
This, of course, is due to the different degrees of dispersion in the 
three colloidal materials. If aluminium hydroxide is to be used, 
by first saturating the cloth with hot aluminium acetate solution 
(p. 560), or by using first aluminium sulphate and then ammonium 
hydroxide, the aluminium hydroxide is precipitated within the 
fibers of the goods. When the material is then dyed, the coloring 
matter is adsorbed by the mordant, with which it forms an in- 
soluble lake, within the fibers. Basic dyes, like Malachite green 
and Methylene blue, behave similarly with tannic acid, or an 
insoluble salt of tannic acid, as mordant. It will be seen that, so 
far as the fabric is concerned, this process, like the first, is a mechan- 
ical one, and is independent of the chemical nature of the goods. 

3. Direct or Substantive Dyes. Most organic dyes are direct 
dyes on silk or wool, and require no mordant with these materials. 
The actions seem to be sometimes chemical, but more often cases 
of adsorption by the silk or wool (both colloids) themselves. A 
few dyes are also fast on cotton. Congo-red Na 2 C32H22N6S 2 06 is 
fast both on cotton and wool, but is no longer much used. 
Chrysophenin is now one of the commonest dyes of this class. 
These dyes, which are sodium salts of complex organic acids, are 
colloids like soap (p. 417), and are salted out within the fibers of 
the goods by adding sodium sulphate to coagulate them and assist 
the adsorption by the cotton. Once adsorbed in this way, unlike 
soap, they cannot be washed out. 

Analytical Reactions of Aluminium Compounds. — The 

alkalies, and alkaline solutions like that of ammonium sulphide, 



566 COLLEGE CHEMISTRY 

precipitate the white hydroxide. The product is soluble in excess 
of the active alkalies. Soluble carbonates also throw down the 
hydroxide. Aluminium compounds, when heated strongly in the 
flame with cobalt salts, give a blue aluminate of cobalt Co(A10 2 )2. 

Exercises. — 1. What are the differences between zinc and 
aluminium, and their corresponding compounds? 

2. Construct equations showing, (a) the hydrolysis of aluminium 
sulphate (p. 558), (6) the interaction of aluminium sulphate and 
cobalt nitrate in the Bunsen flame. 

3. Formulate the ionization of aluminium hydroxide (pp. 557, 
531). 

4. Why does zinc hydroxide, in spite of its feebleness as a base, 
dissolve in ammonium hydroxide, while aluminium hydroxide does 
not? 



CHAPTER XLI 

GERMANIUM, TIN, LEAD 

The metallic elements of the fifth column of the periodic table 
are germanium (Ge, at. wt. 72.5), tin (Sn, at. wt. 119), and lead 
(Pb, at. wt. 207.2). These are on the right side, while titanium 
(Ti, at. wt. 48.1), zirconium (Zr, at. wt. 90.6), cerium (Ce, at. wt. 
140.25), and thorium (Th, at. wt. 232.4) occupy the left side. 

The Chemical Relations of the Family. — All of these ele- 
ments show a maximum valence of four. Germanium, tin, and 
lead are also bivalent. In this respect they resemble carbon and 
differ from silicon, which is more closely allied to the elements on 
the left side of the column. The oxides and hydroxides in which 
these three elements are bivalent become more basic, and the 
elements themselves more metallic in chemical relations, with 
increase in atomic weight. In this they resemble the potassium, 
calcium, and gallium families. Curiously enough, the same three 
hydroxides are also acidic. They are more strongly acidic than 
is zinc hydroxide, for the salts they form by interaction with bases 
are less hydrolyzed than are the zincates. This acidic character 
likewise increases in the order in which the elements are named 
above. 

Germanium 

Germanium (p. 301) forms two oxides GeO and Ge0 2 correspond- 
ing to those of carbon and of tin. Germanious oxide is not very 
definitely basic or acidic, and the sulphide is the only other well- 
defined compound of this set. Germanic oxide and hydroxide are 
acidic entirely. The resemblance to carbon is shown in the for- 
mation of an unstable compound with hydrogen, of germanium 
chloroform GeHCl3 and of a volatile chloride GeCU (b.-p. 87°). 

Tin 

The Chemical Relations of the Element. — Tin is both biva- 
lent and quadrivalent. Each of the oxides and hydroxides SnO 

567 



568 COLLEGE CHEMISTRY 

and Sn(OH) 2 , Sn0 2 and SnO(OH) 2 (or Sn(OH) 4 ), is both basic and 
acidic, so that there are really four series of compounds. Still, 
stannous hydroxide is mainly a base, of a feeble sort, while stannic 
hydroxide is mainly an acid. Thus we have stannous chloride, 
sulphate, and nitrate, which are stable, although they are all more 
or less hydrolyzed by water, and sodium stannite Na 2 .Sn0 2 which 
is unstable. On the other hand, stannic nitrate, sulphate, and 
chloride are completely hydrolyzed by water, while sodium stan- 
nate Na 2 Sn0 3 is comparatively stable. The dioxide Sn0 2 is an 
infusible solid, resembling silicon dioxide. Tin has a tendency 
to give complex acids and salts, like H 2 SnCl 6 , (NH 4 ) 2 .SnCl6, but 
these are ionized also to a small extent after the manner of double 
salts, giving ions of Sn ++++ . Tin forms no salts with weak acids, 
like carbonic acid. 

Occurrence and Extraction. — Tin has long been in use, 
specimens of it being found in Egyptian tombs. The chief ore of 
tin is tinstone, or cassiterite Sn0 2 , which forms square-prismatic 
crystals whose dark color is due to the presence of iron compounds. 
The ore is roughly pulverized and washed, to remove granite or 
slate with which it is mixed, and is then roasted, to oxidize the 
sulphides of iron and copper, and drive off the arsenic which it 
contains. After renewed washing to eliminate sulphate of copper 
and oxide of iron, it is reduced with coal in a reverberatory furnace. 
The tin is afterwards remelted at a gentle heat, and the pure metal 
flows away from compounds of iron and arsenic. The metal is 
produced mainly in Banca and other parts of the East Indies, in 
Bolivia, and in Cornwall. 

Physical and Chemical Properties. — Tin is a silver-white, 
crystalline metal of low tenacity but great malleability (tin-foil). 
Its specific gravity is 7.3, and its melting-point about 232°. 

Tin is dimorphous (p. 266). In 1851, the tin pipes of an organ 
were found to have turned largely into a gray powder. In 1868 a 
shipment of blocks of tin stored in the custom house in Petrograd 
was found to have changed in the same way. Objects of tin in 
museums frequently show spots indicating the presence of the 
"tin pest," as it was called. It now appears that white, metallic 
tin is stable only above 18°, and that below this temperature it is 



tin 569 

unstable and is liable to change into gray tin. This transition 
point is similar to that of sulphur at 96° (p. 265). By immersing 
the tin in a solution of pink-salt (see below), the change is hastened. 
When the two kinds of tin are used as the poles of a cell, and are 
surrounded by pink-salt solution, no difference in potential is 
observed at 18°. But below 18°, white tin, being unstable, is 
more active and becomes negative {^giving positive ions), while 
above 18°, gray tin becomes negative. 

Tin-plate is made by dipping carefully cleaned sheets of mild 
steel into molten tin. Vessels of copper are also coated, internally, 
with tin, to prevent the formation of the basic carbonate (p. 503). 
For this purpose they are cleaned with ammonium chloride, 
sprinkled with rosin (to reduce the oxide), and heated to 230°. 
Molten tin is then spread on the surface with a piece of tow. 
Alloys of tin, such as bronze (p. 503), soft solder (50 per cent lead), 
pewter (25 per cent lead), and britannia metal (10 per cent anti- 
mony and some copper), are much used in the arts. On account 
of the action of soft water containing dissolved oxygen on lead 
(see p. 574), tin pipes are preferred for distributing distilled water 
and for beer pumps. 

Much tin is now recovered by treating old "tin cans" and scrap 
tin-plate with dry chlorine. The dried gas converts the tin into 
stannic chloride SnCU, which is used to make mordants, but hardly 
attacks the iron (p. 160). The process is called detinning. 

Tin, although it displaces hydrogen from dilute acids, is not 
tarnished by moist air. With warm hydrochloric acid it gives 
stannous chloride SnCl 2 and hydrogen. Hot, concentrated sul- 
phuric acid forms stannous sulphate SnS04 and sulphur dioxide 
(cf. p. 276). Nitric acid, when cold and dilute, interacts with it, 
giving stannous nitrate Sn(N03) 2 , and a portion of the nitric acid 
is reduced to ammonia (cf. p. 354). With concentrated nitric 
acid, stannic nitrate is formed, but most of this salt is hydrolyzed 
by the water at the high temperature of the action (cf. p. 535), 
and metastannic acid (H 2 Sn0 3 )5 (/?-stannic acid) remains. The 
final result is shown by the equation (simplified) : 

Sn + 4HN0 3 -> H 2 Sn0 3 + 4N0 2 + H 2 0. 

Tin also displaces hydrogen from caustic alkalies, giving a meta- 
stannate, such as sodium metastannate Na^nC^. 



570 COLLEGE CHEMISTRY 

Chlorides of Tin. — Stannous chloride SnCl 2 ,2H 2 is made by 
the interaction of tin and hydrochloric acid. When the crystals 
are heated, or when a strong aqueous solution is diluted, the salt is 
partially hydrolyzed. In the latter case the basic chloride 
Sn(OH)Cl is deposited. By presence of excess of hydrochloric 
acid, the hydrolysis is prevented. The solution is used as a mor- 
dant (p. 565). 

Stannous chloride tends to pass into stannic chloride SnCU, and 
is therefore an active reducing agent. Thus, it reduces the 
chlorides of mercury (p. 534) and of the noble metals, liberating 
the free metals. The action is of the form Hg++ + Sn++ — > Hg 
+ Sn ++++ . It also reduces free oxygen, or, what is the same thing, 
is oxidized by the air. In this case, stannic chloride is formed in 
the acid solution and the liquid remains clear; in the neutral 
solution a precipitate of the basic chloride is formed as well: 

6SnCl 2 + 2H 2 + 2 -> 4Sn(OH)Cl + 2SnCl4. 

Powdered tin, if placed with the acid solution, will undo the effects 
of this action by reducing the stannic salt to the stannous condition. 
When chlorine acts upon tin, or upon stannous chloride (either 
solid or dissolved), stannic chloride SnCU is formed. The com- 
pound is a colorless liquid (b.-p. 114°) which fumes very strongly 
in moist air, giving hydrochloric acid and stannic acid. It is 
almost completely hydrolyzed by water. The stannic acid which 
is formed is not precipitated, however, but remains in colloidal 
suspension: 

SnCU + 4H 2 <=± 4HC1 + Sn(OH) 4 . 

The chloride, with small amounts of water, gives hydrates, of 
which SnCl4,5H 2 0, "oxymuriate of tin," is used as a mordant. 
Double (or perhaps complex) salts, such as ammonium-stannic 
chloride or " pink-salt" (NH 4 ) 2 SnCl 6 (used as a mordant on cot- 
ton), are readily formed. 

Stannic bromide SnBr 4 (b.-p. 201°) resembles stannic chloride. 

a-Stannic Acid and its Salts. — When a solution of stannic 
chloride is treated with ammonium hydroxide, a white, gelatinous 
precipitate of a-stannic acid is formed: 

SnCU + 4NH4OH -> 4NH4CI + F 2 Sn0 3 + H 2 0. 



TIN 571 

The precipitate loses water gradually until the dioxide remains, 
and neither Sn(OH) 4 nor SnO(OH) 2 is obtainable as a definite 
compound. When stannic oxide is fused with caustic soda, sodium 
metastannate, or a-stannate Na 2 Sn03,3H 2 0, is formed: 

Sn0 2 + 2NaOH -> Na^SnOs + H 2 0. 

This compound is used as a mordant under the name of " pre- 
paring salt." When its solution is acidified, a-stannic acid, the 
actual mordant, is formed by double decomposition. This a- 
stannic acid interacts readily with acids and alkalies, and the 
chloride obtained from it is identical with stannic chloride de- 
scribed above. 

Flannelette and other cotton goods are rendered non-inflam- 
mable by saturation first with sodium a-stannate solution and then, 
after drying, with ammonium sulphate. The acid is too feeble to 
form an ammonium salt: 

Na2Sn0 3 + (NHO2SO4 -> Na 2 S0 4 + SnO(OH) 2 + 2NH 3 . 

The sodium sulphate is washed out and the goods, after being 
dried, contain stannic oxide. The latter cannot afterwards be 
removed by washing, and the material is permanently fireproof. 
Silk is also loaded with stannic oxide, the amount used varying 
from 25 to 300 per cent or more. 

The a-stannates of the metals, aside from those of potassium and 
sodium, like the silicates and carbonates which they much resemble, 
are all insoluble in water, and may be made by double decompo- 
sition. 

^-Stannic Acid, or Metastannic Acid. — The product of the 
action of nitric acid upon tin (p. 569) is a hydrated stannic oxide 
like the foregoing substance, but is not identical with it. It is not 
easily acted upon by alkalies. By boiling it with caustic soda, 
however, and then extracting with pure water, a soluble sodium 
P-stannate N^SnoOn is obtained. /3-stannic acid is also very 
slowly attacked by acids, and the chloride secured from it is not 
identical with the ordinary chloride. For these reasons it is sup- 
posed to be a hydrate of a polymer of stannic oxide (Sn0 2 ) 5 ,- 
zH 2 0. When fused with caustic soda, it gives the same a-stannate 
as does the dioxide itself. 



572 COLLEGE CHEMISTRY 

The Oxides of Tin. — When stannous oxalate is heated in 
absence of air, stannous oxide SnO remains: SnC 2 4 — >SnO+C0 2 
+ CO. It is a black powder which burns in the air, giving the 
dioxide. The corresponding hydroxide Sn 2 0(OH) 2 is formed by 
adding sodium carbonate to stannous chloride solution. It is a 
white powder, easily dehydrated, and interacts with alkalies to 
give soluble stannites, such as Na 2 Sn0 2 . With acids, the hydrox- 
ide gives stannous salts. 

Stannic oxide Sn0 2 is found in nature (p. 568), and may be made 
in pure form by igniting /3-stannic acid. When heated, it becomes 
yellow, but recovers its whiteness when cooled (c/. Zinc oxide, p. 
528). Prepared at a low temperature, it interacts easily with 
acids, but after strong ignition, is affected by them very slowly. 

The Sulphides of Tin. — Stannous sulphide SnS is obtained as 
a dark-brown precipitate when hydrogen sulphide is led into a 
solution of a stannous salt. 

Stannic sulphide SnS 2 is formed likewise by precipitation, and is 
yellow in color. Stannic sulphide loses sulphur when strongly 
heated, and leaves stannous sulphide. It is not much affected by 
dilute acids, but interacts with solutions of ammonium sulphide 
(or sodium sulphide), giving a soluble complex sulphide, namely, 
ammonium sulphostannate : 

SnS 2 + (NH4) 2 S -> (NH 4 ) 2 .SnS 3 . 

The corresponding sodium sulphostannate is easily crystallized in 
the form Na 2 SnS 3 ,2H 2 0. Stannous sulphide is not affected by 
soluble sulphides, but polysulphides, such as yellow ammonium 
sulphide, give with it the above-mentioned sulphostannates : 

SnS + (NH4) 2 S 2 -> (NH4) 2 .SnS 3 . 

With acids the sulphostannates undergo double decomposition, 
but the free acid H 2 .SnS 3 thus produced is unstable and breaks up, 
giving off hydrogen sulphide, and depositing stannic sulphide. 

Analytical Reactions of Salts of Tin. — The two ionic 
forms of tin, Sn _H - and Sn ++++ , are both colorless. Their behavior 
is different. They give a brown and a yellow sulphide, respec- 
tively, with hydrogen sulphide. These sulphides interact with 



LEAD 573 

yellow ammonium sulphide (above). The reducing power of 
stannous-ion Sn ++ is very characteristic (p. 570). The oxides are 
reduced by charcoal in the reducing part of the Bunsen flame and 
the metal is liberated. 

Lead 

The Chemical Relations of the Element. — Lead is both 
bivalent and quadrivalent. The oxides PbO and Pb02, and the 
corresponding hydrated oxides, are all both basic and acidic. Lead 
monoxide is a fairly active base, comparable with cupric oxide, but 
lead dioxide is a feeble one. Both are feebly acidic. The salts of 
bivalent lead, like Pb(N03) 2 , commonly called the plumbic salts, 
are somewhat hydrolyzed by water, but less so than are those of 
tin. The tetrachloride and other salts of quadrivalent lead are 
completely hydrolyzed. The plumbites Na 2 .Pb0 2 and plumbates 
Na 2 .Pb03, like the stannites and stannates, are hydrolyzed to a 
considerable extent. All the compounds in which lead is quad- 
rivalent give up half of the negative radical readily, and are re- 
duced to the " plumbic " condition. The metal displaces hydrogen 
with difficulty, and is easily displaced by zinc. Lead compounds 
are all poisonous, and the effects of repeated, very minute doses 
are cumulative, — resulting in "lead colic. " For this reason, the 
manufacture of white lead is forbidden by law in France, and is 
subject to strict regulation in other countries. 

Occurrence and Metallurgy. — Commercial lead is almost all 
obtained from galena PbS, which crystallizes in cubes, and is found 
in the United States (one-third of the world's supply), Spain, and 
Mexico. This ore often contains considerable amounts of silver 
sulphide AgsS (cf. p. 513). 

The sulphide of lead is first roasted until a sufficient proportion 
of it has been converted into the oxide and sulphate. The furnace- 
doors are then closed, and the temperature raised in order that 
these products may interact with the unchanged part of the 
sulphide: pbg + 2pb0 _^ 3pb + g0 ^ 

PbS + PbS0 4 -* 2Pb + 2S0 2 . 

Another plan consists in heating galenite with scrap iron or iron 
ores and coal: PbS + Fe — » Pb + FeS. The molten ferrous 
sulphide rises to the top as a matte. 



574 COLLEGE CHEMISTRY 

Lead is refined electrolytically by the Betts process. Heavy 
plates of the crude lead form the anodes, thin sheets of pure lead 
the cathodes, and a solution of lead fluosilicate PbSiF 6 the cell 
liquid. The operation is similar to that for refining copper (p. 
511). Silver, gold and bismuth are left as a sludge. 

Physical and Chemical Properties. — Metallic lead is gray 
in color, very soft, and of small tensile strength. Its specific 
gravity is 11.4, and its melting-point 327.4°. While warm, it is 
formed by hydraulic pressure into pipes which are used in plumbing 
and for covering electric cables. On account of its very slow inter- 
action with most substances, sheet lead is used in chemical fac- 
tories, for example, to line sulphuric-acid chambers. An alloy 
containing 0.5 per cent of arsenic is used in making small shot and 
shrapnel bullets. Type-metal contains 20-25 per cent of antimony 
and expands on solidifying, giving a perfect reproduction of the 
mold. In both cases greater hardness is secured by the addition 
of the foreign metal. Solder contains 50 per cent of tin and, being 
a solution, melts at a low temperature. 

Lead oxidizes very superficially in the air. The suboxide Pb 2 
is supposed to be first formed. The final covering is a basic car- 
bonate. Contact with hard waters confers upon lead a similar 
coating composed of the carbonate and the sulphate. These de- 
posits, being insoluble and strongly adherent, enclose the metal 
and protect the water from contamination with lead compounds. 
Pure rain-water, however, since it has no hardness, and contains 
oxygen in solution, gives the hydroxide Pb(OH) 2 , which is notice- 
ably soluble. Hence lead pipes can safely be used only with 
somewhat hard water. When heated in the air, lead gives the 
monoxide PbO or minium Pb 3 04, the latter at lower temperatures. 

The metal displaces hydrogen from hydrochloric acid slowly. 
It is hardly affected by cold concentrated sulphuric acid (cf. p. 284). 
Nitric acid attacks it readily, giving lead nitrate and oxides of 
nitrogen (p. 354). 

Chlorides and Iodide. — Plumbic chloride PbCl2 is precipi- 
tated when a soluble chloride is added to a solution of a lead salt. 
It is slightly soluble in water (1.5 : 100) at 18°, and much more so 
at 100°, 



LEAD 575 

Lead tetrachloride PbCU is a solid at — 15°, and loses chlorine at 
the ordinary temperature. It is made by passing chlorine into 
plumbic chloride suspended in hydrochloric acid. The solution 
contains H 2 PbCl6- Ammonium chloride is added and ammonium 
chloroplumbate (NH^PbCle crystallizes out. When this is 
thrown into cold, concentrated sulphuric acid, an oil, PbCU, 
settles to the bottom. The oil fumes in the air, and closely re- 
sembles stannic chloride SnCU. With little water, it slowly de- 
posits PbClo and gives off chlorine. With much water it is quickly 
hydrolyzed, and lead dioxide is thrown down: 

PbCU + 2H 2 -* Pb0 2 + 4HC1. 

The yellow lead iodide Pbl 2 is formed by precipitation. It crys- 
tallizes in yellow scales from solution in hot water. 

Oxides and Hydroxides. — There are five different oxides of 
lead, Pb 2 0, PbO, Pb 3 4 , PD2O3, and Pb0 2 . The suboxide Pb 2 is a 
dark-gray powder, formed by gently heating the oxalate. Plumbic 
oxide, or lead monoxide PbO, is made by cupellation (p. 513) of 
lead, and the solidified, crystalline mass of yellowish-red color is 
sold as litharge. All the other oxides yield this one when they 
are heated above 600° in the air. Plumbic oxide takes up carbon 
dioxide from the air, and therefore usually contains a basic car- 
bonate. The oxide is used in making glass and enamels and for 
preparing salts of lead. Mixed with glycerine, it gives a cement 
for glass or stone. 

Plumbic hydroxide Pb(OH) 2 is formed by precipitation. It 
gives up water in stages, the successive products being Pb(OH) 2 , 
Pb 2 0(OH) 2 , Pb 3 2 (OH) 2 . These substances are equivalent in 
composition to PbO,H 2 0, 2PbO,H 2 0, and 3PbO,H 2 respectively. 
The hydroxide is observably soluble in water, and gives a solution 
with a faintly alkaline reaction. With acids it forms salts of lead. 
It interacts also with potassium and sodium hydroxides to form 
the soluble plumbites, like sodium plumbite Na 2 .Pb0 2 . 

Minium, or red lead, Pb 3 04, gives off oxygen when heated: 
2Pb 3 4 <=* 6PbO + 2 . 

On account of unequal heating during manufacture, commercial 
red lead is never fully oxidized, and always contains litharge. 
Conversely, commercial litharge usually contains a little minium. 



576 COLLEGE CHEMISTRY 

Minium, when heated with warm, dilute nitric acid, is decom- 
posed, and leaves lead dioxide as an insoluble powder. Two- 
thirds of the lead is basic and one-third is acidic. Minium is 
therefore lead orthoplumbate (see below) : 

Pb 2 .Pb0 4 + 4HN0 3 T± 2Pb(N0 3 ) 2 + HiFbO*. 

The double decomposition as a salt that it thus undergoes is fol- 
lowed by dehydration of the plumbic acid, which is unstable 
(ELtPbC^— »Pb0 2 + 2H 2 0), and the dioxide remains. Red lead 
is used in glass-making, and, when mixed with oil, gives a red paint. 
Lead dioxide Pb0 2 may be obtained as described above in the 
form of a brown powder. It is usually made by adding bleaching 
powder to an alkaline solution of plumbic hydroxide: 

Na 2 .Pb0 2 + Ca(OCl)Cl + H 2 -> 2NaOH + CaCl 2 + Pb0 2 |. 

In this action we may regard the free lead hydroxide, formed by 
hydrolysis of the plumbite, as being oxidized by the bleaching 
powder. Lead dioxide is an active oxidizing agent. It interacts 
with, and sets fire to, a stream of hydrogen sulphide, and it liber- 
ates chlorine from hydrochloric acid. With acids it gives no 
hydrogen peroxide, and is not a peroxide (peroxidate) in the re- 
stricted sense of the term (p. 223). Lead dioxide interacts with 
potassium and sodium hydroxides, giving soluble plumbates. The 
potassium salt K 2 Pb0 3 ,3H 2 is analogous to the metast annate 
K 2 Sn03,3H 2 (p. 571). A mixture of calcium carbonate and lead 
monoxide absorbs oxygen when heated in a stream of air, and the 
yellowish-red calcium orthoplumbate is formed: 

4CaC0 3 + 2PbO + 2 <=* 2Ca 2 Pb0 4 + 4C0 2 . 

The action is reversible, and is at the basis of Kassner's method of 
manufacturing oxygen from the air. 

Other Salts of Lead. — Lead nitrate Pb(N0 3 ) 2 may be made 
by treating lead, lead monoxide, or lead carbonate with nitric acid. 
It forms white, anhydrous octahedra. The nitrate and acetate 
(see below) are the salts of lead which, because of their solubility 
(see Table), are most commonly used. On account of hydrolysis, 
the solution of the nitrate is acid in reaction. 

Lead carbonate PbC0 3 is found in nature. It may be formed as 
a precipitate by adding sodium bicarbonate to lead nitrate solution. 



LEAD 577 

With normal sodium carbonate, a basic carbonate Pb 3 (OH) 2 (C03)2 
is deposited. This basic salt is identical with white lead, which on 
account of its superior opacity, has better covering power than 
zinc-white (p. 528) or permanent white (p. 496). The substance 
is manufactured in various ways, all of which involve the oxidation 
of the lead by the air, the formation of a basic acetate by the inter- 
action of vinegar or acetic acid with the oxide, and the subsequent 
decomposition of the salt by carbon dioxide. The best quality is 
obtained by the Dutch method. In this, gratings of cast lead 
(" buckles") are placed above a shallow layer of vinegar in small 
pots. These pots are buried in manure, which by its decomposition 
furnishes the carbon dioxide and the necessary warmth. The grat- 
ings are gradually converted into a white mass of the basic car- 
bonate. The vapor of acetic acid arising from the vinegar may 
be regarded as a catalytic agent, since it is used over and over 
again. 

Lead acetate Pb(C2H 3 02)2,3H 2 is made by the'action of acetic 
acid on litharge. It is easily soluble in water and, from the sweet 
taste of the solution, is named sugar of lead (used in medicine). 
The basic salt Pb(OH)(C 2 H 3 02) is formed by boiling a solution of 
lead acetate with excess of litharge. Unlike most basic salts, this 
basic salt is soluble in water, and its solution has a faintly alkaline 
reaction. 

Lead sulphate PbSC>4 occurs in nature as anglesite. Being insol- 
uble in water, it is easily obtained by precipitation. 

Natural lead sulphide PbS (galena) forms black, cubic crystals 
with a silvery luster. The precipitated salt is amorphous. It is 
more easily attacked by active acids than is mercuric sulphide (cf. 
p. 531). 

The Storage Battery, — In the ordinary lead accumulator the 

plates consist of leaden gratings. The openings are filled with 
finely divided lead in one plate and with lead dioxide in the other. 
These, and the dilute sulphuric acid in the cell, are the active sub- 
stances when the cell is charged. When the battery is used, the 
S0 4 = ions migrate towards the plates filled with the lead (Fig. 131), 
and convert this into a mass of the insoluble lead sulphate: 
SC>4 = + Pb — > PbS0 4 + 20. These plates receive the negative 
charges. Simultaneously, the H + ions move towards the other 



578 



COLLEGE CHEMISTRY 



plates and there reduce to monoxide the lead dioxide with which 
they are filled. 

Pb0 2 + 2H+ -> H 2 + PbO + 20 . 

These plates acquire positive charges and, by interaction of the 
lead monoxide with the sulphuric acid, become filled, like the 
negative plates, with lead sulphate. During the discharge, much 
sulphuric acid is thus removed from the cell fluid, and the approach- 
ing exhaustion of the cells can thus be ascertained by measuring 
the specific gravity of the fluid. The E.M.F. of the current is a 
little over 2 volts. 



PbS0 4 




Fig. 131. 



Fig. 132. 



The charging is done by passing a current through the cell, in 
the opposite direction to the one which it yields (Fig. 132). The 
H + ions are attracted to the negative plate and an equivalent 
number of S0 4 — ions are formed, so that only lead remains: 

PbS0 4 + 2H+ + 20 -> Pb + 2H+ + S0 4 =. 

Simultaneously, the S0 4 = is attracted by the positive plate and, 
with the lead sulphate there present, forms lead persulphate: 
S0 4 = + PbS0 4 + 2© -> Pb(S0 4 ) 2 . The persulphate, being a 
salt of quadrivalent lead, is at once hydrolyzed and the filling of 
this plate is thus changed into lead dioxide: Pb(S0 4 ) 2 + 2H 2 — > 
Pb0 2 + 2H 2 S0 4 . Both plates are thus brought back to the con- 
dition in which they were before the discharge. 



LEAD 579 

The last set of charges consumes energy, while the first set 
liberates energy. Both may be stated in a single equation: 

charge — * 
2PbS0 4 + 2H 2 <± Pb + 2H 2 S0 4 + Pb0 2 . 
«— discharge 

In the Edison cell, when charged, one plate is of iron and the 
other contains nickelic oxide Ni 2 3 . The cell liquid is a solution 
of potassium hydroxide. When the cell operates, the nickelic 
oxide is reduced to Ni(OH) 2 and the iron is oxidized to Fe(OH) 2 , 
an action which delivers energy: 

Fe + 3H 2 + Ni 2 3 *± Fe(OH) 2 + 2Ni(OH) 2 / 

When the cell is charged, the nickel is reoxidized and the iron 
reduced. 

Paints. — A paint usually contains three ingredients: 

1. The oil hardens to a tough resin on exposure to the air 
(" dries") and adheres firmly to the surface being painted. 

2. The body is a fine powder which makes the paint opaque. 
Since the powder does not shrink, it also "fills" the paint and pre- 
vents the formation of minute pores which otherwise would appear 
in the oil after drying. White lead (p. 577) is a common material 
for the body, but zinc oxide, lithopone (p. 497) and other sub- 
stances are used. 

3. Except in the case of white paint, a pigment is added. Vari- 
ous oxides, such as minium, colored salts, and lakes (p. 565) are 
used as coloring matters. 

The oil does not "dry" by evaporation but gives a resin by 
oxidation. Linseed oil and hemp oil are commonly used. They 
contain glyceryl esters (p. 414) of unsaturated acids, such as that 
of linoleic acid C 3 H 5 (C0 2 Ci7H 3 i) 3 , which contains four units of 
hydrogen less than stearic acid. The unsaturated part of the 
molecule takes up the oxygen. By previously boiling the oil with 
manganese dioxide and other oxides, it is rendered more active, 
and "dries" more quickly. 

Plumbers use a cement made of minium and linseed oil, in which 
the former oxidizes the latter, without access of air being necessary. 



580 COLLEGE CHEMISTRY 

Analytical Reactions of Lead Compounds. — Hydrogen 
sulphide precipitates the black sulphide, even when dilute acids 
are present. Sulphuric acid throws down the sulphate. Potas- 
sium hydroxide gives the white hydroxide, which dissolves in 
excess to form the plumbite. Potassium chromate or dichr ornate 
(q.v.) gives a yellow precipitate of lead chromate PbCr04, which is 
used as a pigment under the name of " chrome-yellow. " 

Titanium, Zirconium, Cerium, Thorium 

The metals on the left side of the fifth column of the periodic 
table are all quadrivalent, although compounds in which a lower 
valence appears are numerous in this family. The first two are 
feebly base-forming as well as feebly acid-forming; the last two 
are base-forming exclusively. 

Titanium occurs in rutile TiTi0 4 . Derived from it are a number 
of titanates of the form K 2 Ti03. Zirconium is found in zircon, the 
orthosilicate of zirconium ZrSi0 4 . The oxide is used in making the 
incandescent substance in some forms of gas lamps. 

Cerium occurs chiefly in cerite [Ce, La, Nd, Pr] Si04,H 2 (cf. p. 
443). The particles of an alloy of cerium (70 per cent) and iron 
(30 per cent), when torn off by a file, catch fire in the air. This 
fact is utilized in making gas-lighters and cigar-lighters. Thorium 
is found in thorite ThSi04, but most of the supply comes from 
monazite sand. The nitrate Th(N0 3 )4,6H 2 is used in making 
Welsbach incandescent mantles (cf. Flame, p. 397). The com- 
pounds are radioactive (see Radium). 

The foundation of the Welsbach mantle is woven of ramie. This 
is saturated with a solution of thorium and cerium nitrates in the 
proportion 99 : 1, and is then molded to the proper shape and dried. 
By heating in a Bunsen flame, the organic matter is burned, and 
the nitrates are decomposed: 

Th(N0 3 ) 4 -> Th0 2 + 4 N0 2 + 2 . 

The oxides retain the form of the fabric and, to prevent breakage 
in handling, the structure is dipped in collodion and dried. 

Exercises, — 1. In what order should you place the elements 
dealt with in this chapter, beginning with the least metallic, and 
ending with the most metallic (p. 436)? 



TITANIUM, ZIRCONIUM, CERIUM, THORIUM 581 

2. Construct equations showing, (a) the interaction of tin and 
concentrated sulphuric acid, (6) of water and. stannous chloride, (c) 
of oxj'gen and stannous chloride in acid solution, (d) the decom- 
position of lead oxalate (p. 575), (e) the interaction of lead mon- 
oxide and acetic acid, (/) and of lead monoxide and lead acetate. 

3. To which class of ionic actions (pp. 259, 270, 504) do the 
reductions by stannous chloride and by tin (p. 570) belong? 

4. What interactions probably occur when lead dioxide liberates 
chlorine from hydrochloric acid? 

5. How should you set about preparing, (a) lead oxalate (in- 
soluble), (6) lead chlorate (soluble)? 

6. Construct equations for the formation of white lead by the 
Dutch process, showing, (1) the formation of the basic acetate by 
the action of oxygen, water, and acetic acid vapor, and (2) the 
action of carbonic acid on the product. 



CHAPTER XLII 

ARSENIC, ANTIMONY, BISMUTH 

jThis family is very closely related to the elements phosphorus 
and nitrogen which precede it in the same column of the periodic 
table. In reading this chapter, therefore, constant reference 
should be made to the chemistry of the corresponding compounds 
of phosphorus. For a general comparison of the elements arsenic 
(As, at. wt. 75), antimony (Sb, at. wt. 120.2) and bismuth (Bi, at. 
wt. 208) with each other and with the two already disposed of, see 
p. 592. It is sufficient here to say that arsenic is mainly an acid- 
forming element, and is therefore a non-metal, while antimony is 
both acid-forming and base-forming, and bismuth is base-forming. 
Each of the three elements gives two sets of compounds, in which 
it is trivalent, and quinquivalent, respectively. None of the 
elements when free displaces hydrogen from dilute acids. 

Arsenic As 

The Chemical Relations of the Element, — Arsenic forms a 
compound with hydrogen AsH 3 . It gives several halogen deriva- 
tives of the type AsX 3 , which are completely hydrolyzed by water. 
Its oxides and hydroxides are acidic. 

Sulphates, nitrates, carbonates, and other salts of arsenic are not 
formed. The complex sulphides (p. 572) are important. 

Occurrence and Preparation. — Arsenic is found free in 
nature. It occurs also in combination with many metals, par- 
ticularly in arsenical pyrites (mispickel) FeAsS. Two sulphides 
of arsenic, orpiment As 2 S 3 and realgar As 2 S 2 , and an oxide As 2 3 , 
are less common. 

The element is obtained either from the native material or by 
heating arsenical pyrites : FeAsS — > FeS + As. During the roast- 
ing of the sulphur ores of metals, arsenic trioxide is formed by the 

582 



ARSENIC 583 

oxidation of the arsenic so frequently present, and collects as a dust 
in the flues. The supply is greatly in excess of the demand. 

Physical and Chemical Properties. — The free element is 
steel-gray in color, metallic in appearance, and crystalline in form. 
It gives off vapor at 180°, and above 600° acquires a vapor pressure 
of 760 mm. The density of the vapor measured at 644° gives 
308.4 as the weight of the G.M.V. (22.4 liters at 0° and 760 mm.). 
The weight of arsenic combining with one chemical unit weight 
(35.46 g.) of chlorine is 25 g. Three times this amount, or 75 g., 
is the smallest weight found in the G.M.V. of any volatile com- 
pound of arsenic, and is therefore accepted as the atomic weight 
(p. 106). Since 308.4 is equal approximately to 4 X 75 (= 300), 
the formula of the vapor of the simple substance at 644° is As 4 . At 
1700° the formula is As 2 (c/. p. 117). 

The free element burns in the air, producing clouds of the solid 
trioxide AS2O3. It unites directly with the halogens, with sulphur, 
and with many of the metals. When boiled with nitric acid, 
chlorine- water, and other powerful oxidizing agents (p. 157), it is 
oxidized in the same way as is phosphorus, and yields arsenic acid 
H 3 As0 4 . 

Arsine AsH 3 . — This substance corresponds in composition to 
ammonia and phosphine, and some of the ways in which it may be 
formed are analogous to those used in the case of these substances. 
Thus, when arsenic and zinc are melted together in the proportions 
to form zinc arsenide Zn 3 As 2 , and the product is treated with dilute 
hydrochloric acid, the result is similar to the action of water or dilute 
acids upon calcium phosphide, and arsine is evolved as a gas: 

Zn 3 As 2 + 6HC1 -» 2AsH 3 + 3ZnCl 2 . 

Arsine (arsenuretted hydrogen) is formed also by the action of 
nascent hydrogen (c/. p. 360) upon soluble compounds of arsenic. 
When a solution of arsenious chloride ASCI3 or arsenic acid is added 
to zinc and hydrochloric acid in a generating flask, arsine is formed: 

AsCl 3 + 3H 2 -> AsH 3 + 3HC1. 

Pure arsine may be secured by leading the mixture with hydrogen 
through a U-tube immersed in liquid air. The arsine (b.-p. —55°) 
condenses as a colorless liquid (m.-p. —119°). 



584 COLLEGE CHEMISTRY 

Arsine burns with a bluish flame, producing water and clouds of 
arsenic trioxide : 2AsH 3 + 30 2 — > 3H 2 + As 2 3 . The combus- 
tion of hydrogen containing arsine, generated as just described, 
gives the same substances. Since arsine, when heated, is readily 
dissociated into its constituents (c/. p. 268), the vapor of free 
arsenic is present in the interior of the hydrogen flame. This 
arsenic may be condensed in the form of a metallic-looking, 
brownish stain by interposition of a cold vessel of white porcelain 
(c/. Fig. 85, p. 268). Even when only a trace of the compound 
of arsenic has been added to the materials in the generator, the 
stain which is produced is very conspicuous. This behavior thus 
furnishes us with the basis of an exceedingly delicate test — 
Marsh's test — for the presence of arsenic in any soluble form of 
combination. The compounds of antimony alone show a similar 
phenomenon (see Stibine). 

Arsine is exceedingly poisonous, the breathing of small amounts 
producing fatal effects. It differs from ammonia more markedly 
than does phosphine, for it is not only without action on water 
or acids, but does not unite directly even with the halides of 
hydrogen. 

Halides of Arsenic. — The halides include a liquid trifluoride 
AsF 3 , a liquid trichloride, a solid tribromide AsBr 3 , and a solid 
tri-iodide Asl 3 . 

The trichloride AsCl 3 , which is prepared by passing chlorine gas 
into a vessel containing arsenic, is easily formed as the result of a 
vigorous action. It is a colorless liquid (b.-p. 130°). When mixed 
with water it is at once converted into the white, almost insoluble 
trioxide. The action is presumably similar to that of water upon 
the corresponding compound of phosphorus (p. 372), but the 
arsenious acid for the most part loses water and forms the insoluble 
anhydride: 

AsCl 3 + 3H 2 <=> As(OH) 3 + 3HCI, 
2As(OH) 3 <=± As 2 3 1 + 3H 2 0. 

This action, however, differs markedly from the other in that it is 
reversible, and arsenic trioxide interacts with aqueous hydrochloric 
acid, giving a solution of arsenious chloride. When this solution 
is boiled, arsenious chloride escapes along with the vapor. 



ARSENIC 585 

Oxides of Arsenic. — Arsenic trioxide AS2O3 is produced by 
burning arsenic in the air and during the roasting of arsenical ores 
(p. 582), and is known as "white arsenic" or simply "arsenic." 
It is purified for commercial purposes by subliming the flue-dust in 
cylindrical pots. The pure trioxide is deposited in a glassy form 
in the upper part of the vessel. Its vapor density shows it to have 
the formula AS4O6. 

When treated with water, the trioxide goes into solution to 
slight extent (1.2 : 100 at 2°), forming arsenious acid, by reversal of 
the second of the actions given above. In boiling water the solu- 
bility is greater (11.5 : 100). When heated in a tube with carbon, 
this oxide is reduced, and the free element, being volatile, is de- 
posited upon the cold part of the tube just above the flame. The 
trioxide is an active poison, since it gradually passes into solution, 
forming arsenious acid. The fatal dose is 0.06-0.18 g. (1-3 grains), 
but "arsenic eaters" become tolerant of it and can take four times 
as much without evil effects. 

The pentoxide AS2O5 is a white crystalline substance, formed by 
heating arsenic acid : 2H 3 As0 4 ,H 2 — > As 2 5 + 4H 2 0. When raised 
to a higher temperature, it loses a part of its oxygen, leaving the 
trioxide. In consequence of this instability, it cannot be formed 
by direct union of oxygen with the trioxide, after the manner of 
phosphorus pentoxide. 

Acids of Arsenic. — When elementary arsenic or arsenious 
oxide is treated with concentrated nitric acid, or with chlorine and 
water, orthoarsenic acid H 3 As0 4 is produced. The substance 
crystallizes as a deliquescent white solid 2H 3 As04,H 2 0. Salts of 
this acid, and of pyroarsenic acid H4As 2 7 and metarsenic acid 
HAs0 3 , corresponding to the phosphoric acids (p. 368), are known. 
The two last acids, themselves, however, are not known as such. 
It has been shown by Menzies that, when the hemihydrate of 
orthoarsenic acid is dried at 100°, the only acid obtainable has the 
composition H 3 As 3 Oio( = 5H 2 0,3As20 5 ). When this acid is heated 
more strongly, it loses water, leaving the pentoxide As 2 5 . With 
metaphosphoric acid, the final elimination of all the water by simple 
heating is impossible. The chocolate-brown silver orthoarsenate 
Ag 3 As04 and the white MgNKiAsC^, like the corresponding phos- 
phates, are insoluble in water. 



586 COLLEGE CHEMISTRY 

Arsenious acid H 3 As0 3 , like sulphurous and carbonic acids, loses 
water, and yields the anhydride (arsenic trioxide) when the attempt 
is made to obtain it from the aqueous solution. The potassium and 
sodium arsenites, K3ASO3 and Na 3 As0 3 , are made by treating arsenic 
trioxide with caustic alkalies, and are much hydrolyzed by water. 
The arsenites of the heavy metals are insoluble, and can be made by 
precipitation. Scheele's green is an arsenite of copper CuHAs0 3 . 
In cases of poisoning by white arsenic, freshly precipitated ferric 
hydroxide (or the same compound in colloidal suspension) or mag- 
nesium hydroxide is administered, since by interaction with the 
arsenious acid they form insoluble substances. 

Sulphides of Arsenic, — Arsenic pentasulphide AS2S5 is ob- 
tained as a yellow powder by decomposition of the sulpharsenates 
(see below), and by leading hydrogen sulphide into the solution of 
arsenic acid in concentrated hydrochloric acid which contains 
AsCl 5 . 

Arsenious sulphide As 2 S 3 occurs in nature as orpiment, and was 
formerly used as a yellow pigment (auripigmentum) . The word 
arsenic is derived from the Greek name for this mineral (dpo-ei/tKov). 
It is obtained as a citron-yellow precipitate when hydrogen sul- 
phide is led into an aqueous solution of arsenious chloride. 

When hydrogen sulphide is led into an aqueous solution of 
arsenious acid, the sulphide is formed, but remains in colloidal 
suspension. It is a negatively charged colloid (p. 417), a small 
amount of H + ion in the liquid rendering the whole electrically 
neutral. It is coagulated by adding solutions of salts, lower con- 
centrations being sufficient the higher the valence of the positive 
ion of the salt (0.05 Molar KC1, 0.0007 M BaCl 2 , 0.00009 M A1C1 3 ). 

Realgar As 2 S 2 is a natural sulphide of orange-red color, and is also 
manufactured by subliming a mixture of arsenical pyrites and 
pyrite: 

2Fe AsS + 2FeS 2 -> 4FeS + As 2 S 2 1 . 

It burns in oxygen, forming arsenious oxide and sulphur dioxide, 
and is mixed with potassium nitrate and sulphur to make " Bengal 
lights." 

Sulpharsenites and Sulpharsenates. — The sulphides of 
arsenic interact with solutions of alkali sulphides after the manner 



ANTIMONY 587 

of the sulphides of tin (p. 572), giving soluble, complex sulphides. 
Arsenious sulphide with colorless ammonium sulphide gives ammo- 
nium sulpharsenite, and with the yellow sulphide gives ammonium 
sulpharsenate : 

3(NH 4 ) 2 S + As 2 S 3 -* 2(NH 4 ) 3 .AsS 3 , 
3(NH4) 2 S + As 2 S 3 + 2S -> 2(NH 4 ) 3 .AsS 4 . 

Proustite (p. 512) is a natural sulpharsenite of silver. 

These salts are decomposed by acids, and give the feebly ionized 
sulpharsenious or sulpharsenic acid: 

(NH4) 3 .AsS 3 + 3HC1 -> 3NH4CI + H 3 AsS 3 -> 3H 2 S t + As 2 S 3 j, 
(NH^.AsS, + 3HC1 -* 3NH 4 C1 + H 3 AsS 4 -» 3H 2 S t + As 2 S 5 i . 

These sulpho-acids, however, at once break up, giving hydrogen 
sulphide as a gas, and the sulphides of arsenic as yellow precipitates. 

Antimony Sb 

The Chemical Relations of the Element. — Antimony 
resembles arsenic in forming a hydride SbH 3 and halides of the 
forms SbXs and SbX 5 . The latter are partially hydrolyzed by 
water with ease, but complete hydrolysis is difficult to accomplish 
with cold water. The oxide Sb 2 3 is basic and also feebly acidic 
(amphoteric), and the oxide Sb 2 05 is acidic. The compositions of 
the compounds are similar to those of the compounds of arsenic, 
but there are in addition salts, such as Sb 2 (S0 4 ) 3 , derived from the 
oxide Sb20 3 . The element gives complex sulphides. 

Occurrence and Preparation. — Antimony occurs free in 
nature. The black trisulphide Sb 2 S 3 , stibnite, is found in Hungary 
and Japan, and forms shining, prismatic crystals. Stibnite is 
roasted in the air in order to remove the sulphur, and the white 
oxide which remains is mixed with carbon and reduced by strong 
heat: 

Sb 2 S 3 + 50 2 -> Sb 2 4 + 3S0 2 , 
Sb 2 4 + 4C-*2Sb + 4CO. 

Properties. — Antimony is a white, crystalline metal, melting 
at 630° (b.-p. 1300°) . It is brittle, and easily powdered. Its vapor 
at 1640° has the formula Sb 2 , while at lower temperatures Sb 4 is 



588 COLLEGE CHEMISTRY 

present. It is used in making alloys such as type-metal, stereo- 
type-metal, and britannia metal (q.v.). The alloys of antimony 
expand during solidification, and therefore give exceptionally sharp 
castings. 

Babbitt's Metal (Sb 3, Zn 69, As 4, Pb 5, Sn 19), and other anti- 
friction alloys used in lining bearings, contain antimony along with 
zinc, copper, and other metals. Molten mixtures of metals 
(alloys), when solidifying, do not always form a homogeneous, 
solid mass. In an anti-friction alloy, what is wanted is a mass, in 
general soft, but containing hard particles. The latter bear most 
of the pressure, yet, as the alloy wears, they are pressed into the 
softer matrix so that a smooth surface is always presented. An 
alloy which has the opposite composition, that is, which gives a 
hard mass containing softer particles, develops heat by friction 
much more rapidly. 

The element unites directly with the halogens. It does not rust, 
but when heated it burns in the air, forming the trioxide Sb 2 03 or a 
higher oxide Sb 2 04. When heated with nitric acid, it yields the 
trioxide and, with more difficulty, antimonic acid H 3 Sb0 4 . 

Stibine ShH 3 . — The hydride of antimony SbH 3 is formed by 
the action of zinc and hydrochloric acid on any soluble compound 
of antimony. By the action of dilute, cold hydrochloric acid on 
an alloy of antimony and magnesium (1 : 2), a mixture of hydrogen 
and stibine containing as much as 11.5 per cent (by volume) of the 
latter may be made. It is separated by cooling with liquid air 
(b.-p. —17°, m.-p. —88°). It is more easily dissociated than is 
arsine (p. 584), and forms a deposit of antimony when a porcelain 
vessel is held in the flame. 

Antimony Halides. — The halides include the trichloride; the 
pentachloride SbCl 5 , a liquid (m.-p. —6°, b.-p. 140°); the tribro- 
mide SbBr 3 , tri-iodide Sbl 3 , trifluoride SbF 3 , and pentafluoride SbFs. 

Antimony trichloride SbCl 3 is made by direct union of chlorine 
and antimony. It forms large, soft crystals (m.-p. 73°, b.-p. 223°), 
and used to be named " butter of antimony." When treated with 
little water, it forms a white, opaque, insoluble basic salt, anti- 
mony oxychloride: 

SbCl 3 + H 2 *± SbOCl I + 2HC1. 



ANTIMONY 589 

With a large amount of water, a greater proportion of the chlorine 
is removed, and Sb 4 5 Cl 2 (= 2SbOCl,Sb 2 3 ) remains. With 
boiling water the oxide is finally formed. The action is not com- 
plete as long as hydrochloric acid is present. It may therefore be 
reversed, so that, on addition of hydrochloric acid to the mixture, 
a clear solution of the trichloride is re-formed. If the concentra- 
tion of the acid is once more reduced by dilution with water, the 
oxychloride is again precipitated. 

Oxides of Antimony. — The trioxide Sb 2 3 (vapor density 
gives Sb 4 6 ) is obtained by oxidizing antimony with nitric acid, or 
by combustion of antimony with a limited supply of oxygen. It is 
a white substance, insoluble in water. It is in the main a basic 
oxide, interacting with many acids to form salts of antimony. 
But it interacts also with alkalies, giving soluble antimonites. 
The pentoxide Sb 2 05 is a yellow, amorphous substance, obtained 
b}' heating antimonic acid. It combines only with bases to form 
salts, and is therefore an acid-forming oxide exclusively. The 
tetroxide St^CX is formed by heating antimony or the trioxide in 
excess of oxygen. It is neither acid- nor base-forming. 

Salts of Antimony. — The nitrate Sb(N0 3 )3 and the sulphate 

Sb 2 (S04) 3 are made by the interaction of the trioxide with nitric and 
sulphuric acids. They are hydrolyzed by water, giving basic salts, 
such as (SbO) 2 S0 4 (= Sb 2 2 S0 4 ), which, like SbOCl, are derived 
from the hydroxide SbO(OH). When the trioxide is heated with 
a solution of potassium bitartrate KHC 4 H0 6 , a basic salt K(SbO)- 
C 4 H40 6 ,iH 2 0, known as tartar-emetic, is formed. This is a white, 
crystalline substance which is soluble in water and is used in 
medicine. The univalent group SbO 1 is known as antimonyl, and 
the above mentioned basic compounds are often called antimonyl 
sulphate, etc. 

Antimonic Acid. — By vigorous oxidation of antimony with 
nitric acid, or by decomposing the pentachloride with water, a 
white, insoluble substance of the approximate composition H 3 Sb0 4 
is obtained. This substance interacts with caustic potash and 
passes into solution. But the salts which have been made are 
pyro- and metantimoniates. Thus, when antimony is fused with 



590 COLLEGE CHEMISTRY 

niter, potassium metantimoniate KSb0 3 is formed. When dis- 
solved, this salt takes up water, giving a solution of the acid 
potassium pyroantimoniate: 

2KSb0 3 + H 2 -> K 2 H 2 Sb 2 7 . 

If this is added to a strong solution of a salt of sodium, an acid 
sodium pyroantimoniate is thrown down, Na 2 H 2 Sb 2 C>7. This is 
almost the only somewhat insoluble salt of sodium. 

Sulphides of Antimony, — The trisulphide Sb 2 S 3 is found in 
nature as the black, crystalline stibnite. As precipitated from 
solutions of salts of antimony, the trisulphide is an orange-red 
powder, which, however, after being melted, assumes the appearance 
of stibnite : 

2SbCl 3 + 3H 2 S <=± Sb 2 S 3 1 + 6HC1. 

Antimony trisulphide, like cadmium sulphide (p. 531), cannot be 
precipitated in presence of concentrated hydrochloric acid. 

The pentasulphide Sb 2 S 5 is obtained by the decomposition of 
the sulphantimoniates (see below) . In appearance it resembles the 
trisulphide and, when heated, decomposes into this substance and 
free sulphur. 

The sulphides of antimony behave towards solutions of the 
alkali sulphides as do the sulphides of arsenic (p. 587). The tri- 
sulphide dissolves in colorless ammonium sulphide with difficulty, 
forming an unstable, soluble ammonium sulphantimonite 

Sb 2 S 3 + 3(NH4) 2 S -> 2(NH4) 3 SbS 3 . 

With the pentasulphide or with yellow ammonium sulphide the 
soluble ammonium sulphantimoniate is readily formed: 

Sb 2 S 5 + 3(NH4) 2 S -> 2(NH4) 3 .SbS 4 , 
Sb 2 S 3 + 3(NH4) 2 S + 2S -> 2(NH4) 3 .SbS 4 . 

The most familiar substance of this class is Schlippe's salt Na 3 SbS4, 
9H 2 0. Pyrargyrite (p. 512) is a natural sulphantimonite. 

When acids are added to solutions of sulphantimoniates, the 
sulphantimonic acid which is liberated decomposes, and antimony 
pentasulphide is thrown down (see under Arsenic, p. 587). 



BISMUTH 591 



Bismuth 



The Chemical Relations of the Element. — Bismuth forms 
no compound with hydrogen. Its compounds with the halogens 
are of the form BiX 3 and are hydrolyzed by water giving basic salts. 
The oxide Bi 2 3 is basic, and the oxide Bi 2 5 is not acidic. Bis- 
muth gives a carbonate, nitrate, phosphate, and other salts, in 
which it acts as a trivalent element. It forms no soluble complex 
sulphides. 

Occurrence and Properties. — This element is found free in 
nature, and also as trioxide Bi 2 03 and trisulphide Bi 2 S 3 . It is a 
shining, brittle metal with a reddish tinge (m.-p. 271°). Bismuth 
is one of the few substances (see water) which expand on solidify- 
ing, the crystals being lighter than the liquid at 271°. It is di- 
morphous, with a transition point (p. 86) at 75°. Mixtures of 
bismuth with other metals of low melting-point fuse at lower 
temperatures than do the separate metals. This is a corollary of 
the fact that a solution freezes at a lower temperature than does 
the pure solvent (p. 134). Thus, Wood's metal, containing bis- 
muth (m.-p. 271°) 4 parts, lead (m.-p. 327°) 2 parts, tin (m.-p. 
232°) 1 part, and cadmium (m.-p. 321°) 1 part, melts at 60.5°, 
considerably below the boiling-point of water. Similar alloys 
are used for safety plugs in steam-boilers and automatic 
sprinklers. 

Bismuth does not tarnish, but when heated strongly it burns to 
form the trioxide. With the halogens it forms a fluoride BiF 3 , a 
bromide BiBr 3 , and an iodide Bil 3 . When the metal is treated 
with oxygen acids, or the trioxide with any acids, salts are pro- 
duced. 

Compounds of Bismuth. — In addition to the basic trioxide 
Bi 2 3 , which is a yellow powder obtained by direct oxidation of the 
metal or by ignition of the nitrate, three other oxides are known — 
BiO, Bi 2 4 , and Bi 2 5 . None of these, however, is either acid- 
forming or base-forming. 

The salts of bismuth, when dissolved in water, give insoluble 
basic salts, and the actions are reversible, the basic salts being 
redissolved by addition of an excess of the acid. In the case of the 



592 COLLEGE CHEMISTRY 

chloride BiCl 3 ,H 2 and the nitrate Bi(N0 3 ) 3 ,5H 2 0, the actions 
taking place are: 

Bids + 2H 2 ?± Bi(OH) 2 Cl + 2HC1, 
Bi(N0 3 ) 3 + 2H 2 <=± Bi(OH) 2 N0 3 + 2HN0 3 . 

The former of these products, when dried, loses a molecule of water, 
giving the oxychloride BiOCl. The oxynitrate Bi(OH) 2 N0 3 is 
much used in medicine, for the treatment of some forms of in- 
digestion, under the name of "subnitrate of bismuth." It is often 
contained in face powders. 

The brownish-black trisulphide Bi 2 S 3 may be obtained by direct 
union of the elements, or by precipitation with hydrogen sulphide. 
This sulphide is not affected by solutions of ammonium sulphide or 
of potassium sulphide. It differs, therefore, markedly from the 
sulphides of arsenic and antimony in its behavior. 

The Family as a Whole 

The elements themselves change progressively in physical 
properties as the atomic weight increases. Nitrogen is a gas 
which with sufficient cooling yields a white solid, phosphorus an 
almost white or a red solid, and arsenic, antimony, and bismuth 
are metallic in appearance. The first combines directly with hy- 
drogen, the next three give hydrides indirectly, and the last does 
not unite with hydrogen at all. The hydride of nitrogen combines 
with water to form a base, while the other hydrides show no such 
tendency. Ammonia unites with acids, including those of the 
halogens, to form salts; phosphine with the hydrogen halides only; 
the others do not combine with acids at all. As regards their 
metallic properties, in the chemical sense, nitrogen and phosphorus 
do not by themselves form positive ions, and furnish us therefore 
with no salts whatever. Arsenic gives a trivalent positive ion, 
which is found in solutions of the halides only. It forms no normal 
sulphates, nitrates, or other salts. Antimony and bismuth both 
give trivalent positive ions. The sulphates, nitrates, etc., of 
antimony, however, are readily decomposed by water with pre- 
cipitation of the hydroxide. The salts of bismuth, on the other 
hand, do not readily give the pure hydroxide with water, although 
they are easily hydrolyzed to basic salts. 



VANADIUM, COLUMBIUM, TANTALUM 593 

The halogen compounds of nitrogen and phosphorus are com- 
pletely hy drolyzed by water, and do not persist when any water is 
present, even when excess of the halogen acid is used. The halogen 
compounds of arsenic are completely hydrolyzed by cold water, but 
exist in solution in presence of excess of the acids. The halogen 
compounds of antimony and bismuth are incompletely hydrolyzed 
by cold water. 

Each element gives a trioxide and a pentoxide. With nitrogen 
these are acid-forming, being the anhydrides of nitrous and nitric 
acids. With phosphorus the trioxide and the pentoxide are an- 
hydrides of acids. With arsenic the trioxide is basic towards the 
halogen acids, and is the first example of a basic oxide which we 
encounter in this group. The pentoxide, however, is acid-forming. 
The trioxide of antimony is mainly base-forming, although it is 
feebly acid-forming also. The pentoxide is acid-forming. The 
trioxide of bismuth is base-forming exclusively, and the pentoxide 
has no derivatives. 

These statements, which could easily be expanded, are sufficient 
to show that when the periodic law is borne in mind it furnishes 
valuable aid in systematizing the chemistry of a group like this. 

Analytical Reactions of Arsenic, Antimony, and Bismuth. 

— The ions which are most frequently encountered are As +++ , 
Sb+++, Bi+++, As0 4 =-, and As0 3 =- The first three, with hydro- 
gen sulphide, give colored sulphides which are not affected by 
dilute acids. The sulphides of arsenic and antimony are separable 
from the sulphide of bismuth by solution in yellow ammonium 
sulphide. Marsh's test enables us to recognize the presence of 
traces of compounds of arsenic and antimony. Oxygen compounds 
of arsenic, when heated with carbon, give a volatile, metallic- 
looking deposit of arsenic. 

Vanadium, Columbium, Tantalum 

Of these elements, vanadium is less uncommon than the others. 
It is found in rather complex compounds. When these are heated 
with soda and sodium nitrate, sodium metavanadate Na VO3 is 
formed, and can be extracted with water. The element forms 
several chlorides, such as VC1 2 , VC1 3 , VCI4, VOCl 3 , and five oxides, 
V 2 0, VO, V2O3, V0 2 , and V 2 5 . The element has very feeble base- 



594 COLLEGE CHEMISTRY 

forming properties, and gives only a few unstable salts. Ferro- 
vanadium, an alloy, is used in making vanadium steel. 

Columbium (or niobium), first discovered and named by Hat- 
chett (1801), and tantalum possess feeble base-forming properties, 
their chief compounds being the columbates and tantalates. 

Exercises. — 1. How do you account for the fact that the 
molecular weight of arsenic at 644° is not exactly 300, and why is 
308.4 -T- 4 not accepted as the atomic weight? 

2. Formulate the series of changes involved in the solution of 
arsenic trioxide and the interaction of hydrochloric acid with the 
arsenious acid so formed (c/. p. 272). 

3. What is the full significance of the fact that arsenic penta- 
sulphide may be precipitated by hydrogen sulphide from a solution 
of arsenic acid in hydrochloric acid? Make the equation. 

4. To what classes of chemical changes do the interactions of 
arsenious sulphide and antimony trisulphide with yellow ammo- 
nium sulphide belong? 

5. Construct equations showing the interaction of, (a) oxygen 
and arsenical pyrites, (b) chlorine-water and arsenic, (c) the de- 
hydration of orthoarsenic acid, (d) potassium hydroxide and arsenic 
trioxide, (e) concentrated nitric acid and antimony, (/) potassium 
bitartrate and antimony trioxide, (g) acids and ammonium ortho- 
sulphantimoniate. 

6. How should you set about making Schlippe's salt? 



CHAPTER XLIII 

THE CHROMIUM FAMILY. RADIUM 

The chromium (Cr, at. wt. 52) family includes molybdenum (Mo, 
at. wt. 96), tungsten (W, at. wt. 184), and uranium (U, at. wt. 
238.2), and occupies the seventh column of the periodic table along 
with the sulphur and selenium family. 

The Chemical Relations of the Family. — The features 
which are common to the four elements are also those which 
affiliate them most closely with their neighbors on the right side 
of the column. They yield oxides of the forms CrOs, Mo0 3 , 
~W0 3 , and U0 3 , which, like S0 3 , are acid anhydrides, and show the 
elements to be sexivalent. They give also acids of the form H 2 X0 4 , 
such as chromic acid H 2 Cr0 4 . These acids correspond to sulphuric 
acid, and their salts, for example the chromates, resemble the 
sulphates. 

Aside from the chromates, the first element forms also two basic 
hydroxides Cr(OH) 2 and Cr(OH) 3 , from which the numerous chro- 
mous (Cr 4-1- ) and chromic (Cr 4 ""*" 4 ") salts are derived. Uranium is 
base-forming, as well as acid-forming. Molybdenum and tungsten 
are not base-forming elements. 

Chromium Cr 

The Chemical Relations of the Element. — Chromium gives 
four classes of compounds, and most of them are colored sub- 
stances (Gk. xp'Vs color). The chromates are derived from 
chromic acid H 2 Cr0 4 , which, however, is itself unstable, and leaves 
the anhydride when the solution is evaporated. The oxide and 
hydroxide in which the element is trivalent, namely Cr 2 3 and 
Cr(OH) 3 , are weakly basic and still more weakly acidic. Hence 
we have chromic salts such as CrCl 3 and Cr 2 (S0 4 )3 which are 
somewhat hydrolyzed, but no carbonate, and no sulphide which 
is stable in water. The compounds in which the same hydroxide 

595 



596 COLLEGE CHEMISTRY 

acts as an acid are the chromites, and are derived from the less 
completely hydrated form of the oxide CrO(OH). Potassium 
chromite K.Cr0 2 is more easily hydrolyzed, however, than is 
potassium zincate or potassium aluminate. Finally, the chro- 
mous salts such as CrCl2 and CrS0 4 correspond to chromous 
hydroxide Cr(OH) 2 in which the element is bivalent. This hy- 
droxide is more distinctly basic than is chromic hydroxide, and 
forms a carbonate and sulphide which can be precipitated in 
aqueous solution. 

Occurrence and Isolation. — Chromium is found chiefly in 
ferrous chromite Fe(Cr0 2 ) 2 , which constitutes the mineral chro- 
mite, and in crocoisite PbCr0 4 , which is chromate of lead. It 
was first discovered in the latter mineral by Vauquelin (1797). 
The metal is easily made by reduction of the oxide with aluminium 
filings by Goldschmidt's method (p. 556). 

Physical and Chemical Properties. — Chromium is a white, 
crystalline, very hard metal (m.-p. 1520°). It does not tarnish, 
but when heated it burns in oxygen, giving the green chromic 
oxide Cr 2 3 . It seems to exist in two states, an active and a pas- 
sive one, the relations of which are still somewhat obscure. A 
fragment which has been made by the Goldschmidt method, or 
has been dipped in nitric acid, is passive, and does not displace 
hydrogen from hydrochloric acid. When, however, the specimen 
is warmed with this acid, it begins to interact, and thereafter 
behaves as if it lay between zinc and cadmium in the electro- 
motive series. If left in the air, it slowly becomes inactive again. 

Tin and iron with hydrochloric acid form stannous and ferrous 
chlorides respectively, because the higher chlorides, if present, 
would be reduced by the active hydrogen (p. 360). Here, for 
the same reason, chromous chloride and not chromic chloride is 
formed: 

Cr + 2HCl->CrCl 2 + H 2 , or Cr + 2H+ -> Cr++ + H 2 . 

Chromium is used in making chrome-steel, for armorplate. 
The strange alloys, which, although composed of active metals, 
are not attacked by acids (even boiling nitric acid), usually contain 
chromium (e.g., 60% Cr, 36% Fe, 4% Mo). 



THE CHROMIUM FAMILY 597 

Derivative of Chromic Acid 

Potassium Chr ornate K 2 CrO^. — This and the sodium salt, or 
rather the corresponding dichrornates (see below), are made di- 
rectly from chromite, and form the starting-point in the prepara- 
tion of the other compounds of chromium. The finely powdered 
mineral is mixed with potash and limestone, and roasted. The 
lime is emplo} r ed chiefly to keep the mass porous and accessible 
to the oxygen of the air, the potassium compounds being easily 
fusible : 

4Fe(Cr0 2 ) 2 + 8K 2 C0 3 + 70 2 -> 2Fe 2 3 + 8K 2 Cr0 4 + 8C0 2 . 

The iron is oxidized to ferric oxide, and the chromium passes from 
the state of chromic oxide in the chromite (FeO,Cr 2 3 ) to that of 
chromic anhydride in the potassium chromate (K 2 0,O0 3 ). Thus, 
more insight is given into the nature of the action by the equation: 

4(FeO,Cr 2 3 ) +8(K 2 0,C0 2 ) +70 2 ^2Fe 2 3 +8(K 2 0,Cr0 3 ) +8C0 2 . 

The cinder is treated with hot potassium sulphate solution. This 
interacts with the calcium chromate, which is formed at the same 
time, giving insoluble calcium sulphate: 

CaCr0 4 + K 2 S0 4 *=► CaS0 4 J + K 2 Cr0 4 . 

The whole of the potassium chromate goes into solution. 

Potassium chromate is pale-yellow in color, gives anhydrous, 
rhombic crystals like those of potassium sulphate, and is very 
soluble in water (61 : 100 at 10°). 

Sodium chromate Na 2 CrO 4 ,10H 2 O is made by using sodium car- 
bonate in the process just described. 

The Dichrornates. — When a solution of potassium sulphate is 
mixed with an equivalent amount of sulphuric acid, potassium 
bisulphate is obtainable by evaporation : K 2 S0 4 + H 2 S0 4 — ■» 
2KHS0 4 . The dry acid salt, when heated, loses water (p. 286), 
giving the pyrosulphate (or disulphate) : 2KHS0 4 *± K 2 S 2 7 + 
H 2 0, but the latter, when redissolved, returns to the condition of 
acid sulphate. The second action is instantly reversed in presence 
of water. Now, when an acid is added to a chromate we should 
expect the chromic acid H 2 Cr0 4 , thus liberated, to interact, giving 



598 COLLEGE CHEMISTRY 

an acid chromate (say, KHCr0 4 ) . No acid chromates are known, 
however, and instead of them, pyrochromates or dichromates are 
produced, with elimination of water. In other words, the second 
of the above actions is not appreciably reversible in presence of 
water when chromates are in question : 

K 2 Cr0 4 +H 2 S0 4 ->(H 2 Cr0 4 )+K 2 S0 4 . 

K 2 CrQ 4 ( + H 2 CrQ 4 ) -> K 2 Cr 2 Q 7 + H 2 Q. 

2K 2 Cr0 4 + H 2 S0 4 -> K 2 Cr 2 7 + H 2 + K 2 S0 4 . (1) 

In terms of the ionic hypothesis, S 2 7 = is unstable in water, and 
interacts with the OH~ ion it contains, giving water and sul- 
phate-ion, while Cr 2 7 — is stable in water and is formed from 
the interaction of water and chromate-ion : 

S 2 7 = +20H-^H 2 + 2S0 4 =, 

Cr 2 7 = + 20H~ ±* H 2 + 2Cr0 4 = (2) 

The dichromates of potassium and sodium are made by adding 
sulphuric acid to the crude solution of the chromate obtained from 
chromite (p. 597). They crystallize when the liquid cools, and the 
mother-liquor, containing the potassium sulphate and undeposited 
dichromate, is used for extracting a fresh portion of cinder. As the 
dichromates are much less soluble than the chromates, they crys- 
tallize from less concentrated solutions, and can therefore be ob- 
tained in purer condition. For this reason the extract is always 
treated for dichromate. 

Potassium dichromate K 2 Cr 2 7 (or K 2 Cr0 4 ,Cr0 3 ) crystallizes in 
asymmetric tables of orange-red color. Its solubility in water is 
8 : 100 at 10° and 12.5 : 100 at 20°. Sodium dichromate Na 2 Cr 2 7 , 
2H 2 forms red crystals also, and its solubility is 109 : 100 at 15°. 
This salt is now cheaper than potassium dichromate, and has 
largely displaced the latter for commercial purposes. 

Chemical Properties of the Dichromates. — 1. When con- 
centrated sulphuric acid is added to a strong solution of a dichro- 
mate (or chromate), chromic anhydride Cr0 3 separates in red 
needles : 

Na 2 Cr 2 7 + H 2 S0 4 -> Na 2 S0 4 + H 2 + 2Cr0 3 J, . 

2. Although a dichromate lacks the hydrogen, it is essentially of 
the nature of an acid salt, just as SbOCl lacks hydroxyl, but is 



THE CHROMIUM FAMILY 599 

essentially a basic salt. Hence, when potassium hydroxide is 
added to a solution of potassium dichromate, potassium chromate 
is formed : 

K 2 Cr 2 7 + 2KOH -> 2K 2 Cr0 4 + H 2 0. 

The solution changes from red to yellow, and the chromate is 
obtained by evaporation. In this way the pure alkali chromates 
are made. 

3. By addition of potassium dichromate to a solution of a salt of 
a metal whose chromate is insoluble, the chromate and not the 
dichromate is precipitated. This occurs in consequence of the 
fact that there is always a little hydrogen-ion and Cr0 4 = (equation 
(2), above) in the solution of the dichromate: 

2Ba(N0 3 ) 2 + K 2 Cr 2 7 + H 2 <=± 2BaCr0 4 1 + 2KN0 3 + 2HN0 3 . 

Being essentially an acid salt, the dichromate produces a salt and 
an acid, as any acid salt would do. For example: 

Ba(N0 3 ) 2 + KHSO4 <=* BaS0 4 J + KN0 8 + HN0 3 . 

4. The dichromates of potassium and sodium melt when heated 
and, at a white heat, decompose, giving the chromate, chromic 
oxide, and free oxygen. To make the equation, we note that the 
dichromate, for example K 2 Cr 2 07, may be written as K 2 Cr0 4 ,Cr0 3 , 
and the Cr0 3 , if alone, will decompose thus : 2Cr0 3 — > Cr 2 3 + 30. 
Since the product must contain a multiple of 2 , the equation is: 

4K 2 Cr 2 7 -+4K 2 Cr0 4 + 2Cr 2 3 + 30 2 . 

5. With free acids the dichromates give powerful oxidizing mix- 
tures, in consequence of their tendency to form chromic salts. 
Since the former correspond to the oxide Cr0 3 and the latter to 
Cr 2 3 , the passage from the former to the latter must furnish 30 
for every 2Cr0 3 transformed. In dilute solutions, unless a body 
capable of being oxidized is present, no actual decomposition, 
beyond the liberation of chromic acid,* occurs. When concen- 
trated hydrochloric acid is used, this acid itself suffers oxidation: 

K 2 Cr 2 7 + 8HC1 -> 2KC1 + 2CrCl 3 + 4H 2 (+ 30) . 

(30) + 6HC1->3H 2 + 3C1 2 . 

K 2 Cr 2 7 + 14HC1 -» 2KC1 + 2CrCl 3 + 7H 2 + 3C1 2 . 

* Not shown as a distinct stage in the subsequent equations. 



600 COLLEGE CHEMISTRY 

When sulphuric acid is employed, an oxidizable substance such as 
hydrogen sulphide (c/. p. 270), sulphurous acid, or alcohol must be 
present, if the dichromate is to be reduced: 

K 2 Cr 2 7 + 4H 2 S0 4 -> K 2 S0 4 + Cr 2 (S0 4 ) 3 + 4H 2 0( + 30) (1) 

(30)+ 3H 2 S0 3 ->3H 2 S0 4 (2) 

or (30) + 3C 2 H 5 OH -» 3C 2 H 4 t + 3H 2 (2') 

[alcohol] [aldehyde] 

In each case the usual summation of (1) and (2), with omission of 
the 30, gives the equation for the whole action. When (1) is dis- 
sected, K 2 0,2Cr0 3 giving Cr 2 3 ,3S0 3 + 30 is found to be its essen- 
tial content. In practice, this sort of action is used for the purpose 
of making chromic salts, and for its oxidizing effects, as in the 
preparation of aldehyde and in the dichromate battery. 

Other Uses of Dichromates. — When paper is coated with 
gelatine containing a soluble chromate or dichromate and, after 
being dried, is exposed to light, chromic oxide is formed by reduc- 
tion, and combines with the gelatine. This product will not swell 
up or dissolve in tepid water, as does pure gelatine. This action 
is used in many ways for purposes of artistic reproduction. Thus, 
if the gelatine mixture is made up with lampblack and, after the 
coating has dried, is covered with a negative and exposed to light, 
the parts which were protected from illumination may afterwards 
be washed away, while the carbon print remains. The gelatine 
layer can be transferred to wood or copper before washing. When 
materials of different colors are substituted for the lampblack, 
prints of any desired tint may be made by the same process. 

Sodium dichromate is used, instead of tan-bark, in tanning kid 
and glove leathers. A reducing agent is employed to precipitate 
chromic hydroxide Cr(OH) 3 in the leather. Its use diminishes 
the time required for the process from 8 or 10 months to a few 
hours. The hide is a mixture of colloidal materials, and the hy- 
droxide is adsorbed. 

Insoluble Chromates. — A number of chromates, formed by 
precipitation with a solution of a soluble chromate or dichromate, 
are familiar. Thus, lead chromate PbCr0 4 is used as a yellow 
pigment. By treatment with limewater it gives a basic salt of 
brilliant orange color — chrome-red Pb 2 OCr0 4 . Salts of calcium 



THE CHROMIUM FAMILY 601 

give a yellow, hydrated calcium chromate CaCr04,2H02 analogous 
to gypsum, and, like it, perceptibly soluble in water (0.4 : 100 at 
14°) . Barium chromate BaCr0 4 is also yellow. It interacts with 
active acids to form the dichromate, and passes into solution. It 
is not soluble enough to be attacked by acetic acid. Strontium 
chromate SrCr04, however, is soluble in acetic acid. Silver chro- 
mate Ag2Cr0.i is red, and interacts easily with acids. It will be 
observed that there is a close correspondence between the relative 
solubilities (see Table) of the chromates and the sulphates. 

Chromic Anhydride Cr0 3 . — This oxide is made as described 
above (par. 1, p. 598), and is often called chromic acid. It is 
soluble in water, and combines with the latter to some extent, 
giving dichromic acid H 2 .Cr 2 07. In a solution acidified with an 
active acid it is much used as an oxidizing agent for organic sub- 
stances. It interacts with acids in the same way as do the dichro- 
mates, giving chromic salts and furnishing oxygen to the oxidizable 
body. When heated by itself, it loses oxygen readily, and yields 
the green chromic oxide: 4Cr03 — > 2Cr 2 3 + 30 2 . 

Chromyl Chloride Cr0 2 Cl 2 * — This compound corresponds to 
sulphuryl chloride S0 2 C1 2 , and is made by distilling a dichromate 
with a chloride and concentrated sulphuric acid: 

K 2 Cr 2 7 + 4KC1 + 3H 2 S0 4 -> 2Cr0 2 Cl 2 t + 3K 2 S0 4 + 3H 2 0. 

The hydrochloric acid liberated from the chloride may be supposed 
to interact with chromic acid from the dichromate: 

Cr0 2 (OH) 2 + 2HC1 -> Cr0 2 Cl 2 + 2H 2 0. 

Chromyl chloride is a red liquid, boiling at 1 18°. It fumes strongly 
in moist air, being hj^drolyzed by water. This action is the re- 
verse of that shown in the last equation. The corresponding 
bromine and iodine compounds are unstable, and when a bromide 
or iodide is treated as described above, the halogens are liberated 
by oxidation, and no volatile compound of chromium appears. 
Hence, when an unknown halide is mixed with potassium dichro- 
mate and sulphuric acid, and distilled, and the vapors are caught 
in ammonium hydroxide, the finding of a chromate in the dis- 



602 COLLEGE CHEMISTRY 

tillate demonstrates the existence of a chloride in the original 
substance: 

Cr0 2 Cl 2 + 4NH 4 OH -> (NH 4 ) 2 Cr0 4 + 2NH 4 C1 + 2H 2 0. 

This action is used as a test for the presence of traces of chlorides 
in large amounts of bromides or iodides. 



Chromic and Chromous Compounds 

Chromic Chloride, — A hydrated chloride CrCl 3 ,6H 2 is ob- 
tained by treating the hydroxide Cr(OH) 3 with hydrochloric acid 
and evaporating. When heated, this hydrate is hydrolyzed, and 
chromic oxide remains. The anhydrous chloride CrCl 3 is formed 
by sublimation, as a mass of brilliant, reddish-violet scales, when 
chlorine is led over heated metallic chromium. In this form the 
substance dissolves with extreme slowness, even in boiling water, 
but in presence of a trace of chromous chloride or stannous chloride 
it is easily soluble. The solution is green, as are all solutions of 
chromic salts after they have been boiled, but on standing in the 
cold, bluish crystals of CrCl3,6H 2 are deposited. These give a 
violet solution containing Cr +++ + 3C1~, but boiling reproduces the 
green color. The green material can also be obtained in crystals 
as a hexahydrate, and is therefore isomeric (p. 421) with the 
violet variety. With the green isomer, in cold solution, silver 
nitrate precipitates at first only one-third of the chlorine as silver 
chloride. 

Chromic Hydroxide. — When ammonium hydroxide is added 
to a solution of a chromic salt, a hydrated hydroxide of pale-blue 
color, 2Cr(OH) 3 ,H 2 0, is thrown down. This interacts with acids, 
giving chromic salts. It also dissolves in potassium and sodium 
hydroxides to form green solutions of chromites of the form KCr0 2 . 
When the solutions of the alkali chromites are boiled, the free 
chromic hydroxide, present in consequence of hydrolysis, is con- 
verted into a greenish, less completely hydrated, and less soluble 
variety. This begins to come out as a precipitate, and soon the 
whole action is reversed. Insoluble chromites, such as that of 
iron Fe(Cr0 2 ) 2 , are found in nature. Many of them, like Zn(Cr0 2 ) 3 
and Mg(Cr0 2 ) 2 , may be formed by fusing the oxide of the metal 



THE CHROMIUM FAMILY 603 

with chromic oxide; the action being similar to that used in making 
zincates (p. 529) and aluminates (p. 557). The hydroxide is used 
as a mordant (p. 565) and is the active substance in the chrome- 
tanning process (p. 600). 

Chromic Oxide Cr 2 O s . — This oxide is obtained as a green, 
infusible powder by heating the hydroxide; or, more readily, by 
heating dry ammonium dichromate; or by igniting potassium 
dichromate with sulphur and washing the potassium sulphate out 
of the residue : 

(XH 4 ) 2 Cr 2 7 -> N 2 + 4H 2 + Cr 2 3 , 

K 2 Cr 2 7 + S -> K 2 S0 4 + Cr 2 3 . 

Chromic oxide is not affected by acids, but may be converted into 
the sulphate by fusion with potassium bisulphate. It is used for 
making green paint, and for giving a green tint to glass. When the 
oxide, or any of the chromic salts, is fused with a basic substance 
such as an alkali carbonate, it passes into the form of a chromate, 
absorbing the necessary oxygen from the air. If an alkali nitrate 
or chlorate is added, the oxidation goes on more quickly. The 
alkaline solution of the chromites may be oxidized, for example by 
adding chlorine or bromine, and chromates are formed. 

Chromic Sulphate Cr Q (SO^) 39 15H 2 0. — This salt crystallizes 
in reddish-violet crystals, and may be made by treating the hy- 
droxide with sulphuric acid. When mixed with potassium sul- 
phate, it gives reddish-violet, octahedral crystals of chrome-alum 
(c/. p. 558), K 2 S0 4 ,Cr 2 (S0 4 ) 3 ,24H 2 0. This double salt is most 
easily obtained by reducing potassium dichromate in dilute sul- 
phuric acid by means of sulphurous acid (p. 600), and allowing 
the solution to crystallize. The solution of the crystals, either of 
the pure sulphate or of the alum, is bluish- violet (Cr +++ ), but 
when boiled becomes green. The green compound is formed by 
hydrolysis and is gummy and uncrystallizable. It even yields 
products which do not show the presence either of the Cr +++ or 
the S0 4 = ion. It seems to be formed thus: 

2Cr 2 (S0 4 ) 3 + H 2 <=* Cr 4 0(S0 4 ) 4 .S0 4 + H 2 S0 4 . 

The green materials revert slowly to the violet ones by reversal of 
the above action when the solution remains in the cold, and so 



604 COLLEGE CHEMISTRY 

crystals of the sulphate or of^the alum are obtainable from the 
green solutions. 

Chromous Compounds. — By the interaction of chromium 
with hydrochloric acid, or by reducing chromic chloride in a stream 
of hydrogen, chromous chloride CrCl2 is formed. The anhydrous 
salt is colorless, and its solution is light blue (Cr 4-1- ). Like stan- 
nous chloride, it is very easily oxidized by the air, a solution of it 
containing excess of hydrochloric acid being used in the laboratory 
to absorb oxygen: 

4CrCl 2 + 4HC1 + 2 ->4CrCl 3 + 2H 2 0. 

Chromous hydroxide Cr(OH) 2 is obtained as a yellow precipitate 
when alkalies are added to the chloride. With sulphuric acid it 
gives chromous sulphate CrS0 4 ,7H 2 0, which is one of the vitriols 
(p. 529). 

Analytical Reactions of Chromium Compounds. — The 

chromic salts give the bluish-violet chromic-ion Cr 4-1-1 ", or the green 
complex cations, and may be recognized in solution by their color. 
The chromates and dichromates give the ions Cr0 4 = and Cr 2 07~, 
which are yellow and red respectively. From chromic salts, 
alkalies and ammonium sulphide precipitate the bluish-green 
hydroxide, and carbonates give a basic carbonate which is almost 
completely hydrolyzed to hydroxide. By fusion with sodium 
carbonate and sodium nitrate, they yield a yellow bead containing 
the chromate. The chromates and dichromates are recognized 
by the insoluble chromates which they precipitate, and by their 
oxidizing power when mixed with acids. All compounds of chro- 
mium give a green borax bead containing chromic borate, and this 
bead differs from that given by compounds of copper (c/. p. 510), 
both in tint and in being unreducible. 

Molybdenum, Tungsten, Uranium 

Molybdenum. — This element is found chiefly in wulfenite 
PbMo0 4 and molybdenite MoS 2 , The latter resembles black 
lead (graphite), and its appearance suggested the name of the 
element (Gk. poXvpSaiva, lead). The molybdenite is converted by 



THE CHROMIUM FAMILY 605 

roasting into molybdic anhydride M0O3. When this is treated with 
ammonium hydroxide, or with sodium hydroxide, ammonium 
molybdate (NH 4 ) 2 Mo0 4 or sodium molybdate Na 2 Mo0 4 , 10H 2 O is 
obtained. The metal itself is liberated by reducing the oxide or 
chloride with hydrogen. ' When pure it is a silvery metal and, 
like iron (q.v.), takes up carbon and shows the phenomena of 
tempering. The oxides Mo 2 3 , Mo0 2 , and M0O3 are known, but 
the lower oxides are not basic. The chlorides M03CI6, M0CI3, 
M0CI4, and M0CI5 have been made. The chief use of molybdenum 
compounds in the laboratory is in testing for and estimating phos- 
phoric acid. When a little of a phosphate is added to a solution 
of ammonium molybdate in nitric acid, and the mixture is warmed, 
a copious } T ellow precipitate of a phosphomolybdate of ammonium 
(NH 4 ) 3 P04,llMo03,6H 2 is formed. The compound is soluble 
in excess of phosphoric acid and in alkalies, but not in dilute 
mineral acids. 

Tungsten. — The minerals scheelite CaW0 4 and wolfram 
[Fe,Mn]W0 4 are tungstates of calcium and of iron and manganese, 
respectively. By fusion of wolfram with sodium carbonate and 
extraction with water, sodium tungstate Na 2 W0 4 ,2H 2 is secured. 
It is used as a mordant and for rendering muslin fireproof. Acids 
precipitate tungstic acid H 2 W0 4 ,H 2 from solutions of this salt. 
The element gives the oxides W0 2 and W0 3 , the latter being 
formed by ignition of tungstic acid. The chlorides WC1 2 , WC1 4 , 
WCI5, and WC1 6 are known, the last being formed directly, and 
the others by reduction. 

The metal has important uses, and the annual production is 
greater than the total of all the metals which follow it in the list 
on p. 436. The metal (sp. gr. 19.6) can be liberated by reduction 
of the oxide by hydrogen or by carbon. It has a higher melting 
point (3540°) than any other metal and, on this account, and be- 
cause it is less volatile than carbon, is now used for filaments in 
electric lamps. A carbon filament also requires 3.25 watts per 
candle power while a tungsten filament uses only 1.25 watts per 
1 c. p. The powdered metal obtained by reduction can be pressed 
into wire form and then rolled while strongly heated by an electric 
current until a compact wire is obtained. The metal can also be 
obtained in massive form by reducing the oxide with aluminium, 



606 COLLEGE CHEMISTRY 

provided the crucible and mixture are heated strongly in advance. 
In 1914, in the United States alone, about a hundred million tung- 
sten lamps were manufactured. Shop work has been almost revolu- 
tionized by the use of tungsten steel tools, which can be used at 
high speed and, even when thus heated red hot by friction, retain 
their temper. Tungsten steel contains tungsten (16 to 20%), 
carbon (0.55 to 0.75%), chromium (2.5 to 5%), and vanadium 
(0.35 to 1.5%). 

Uranium. — Pitchblende, which contains the oxide U 3 8 along 
with smaller amounts of many other elements, is found mainly 
in Joachimsthal (Bohemia) and in Cornwall. Carnotite, a ura- 
nate and vanadate of potassium K 2 0,2U03,V 2 05,3H20 occurs in 
Colorado. Pitchblende is roasted with lime, the calcium uranate 
CaU04 thus formed is decomposed with sulphuric acid, giving 
uranyl sulphate UO2SO4. When excess of sodium carbonate is 
added to the solution of the latter, the foreign metals are precipi- 
tated and sodium diuranate Na2U 2 07,7H 2 0, which is also thrown 
down, dissolves in the excess as Na 2 U0 4 . 

After filtration, the diuranate of sodium is reprecipitated by 
neutralizing with sulphuric acid and boiling. This salt is used 
in making uranium glass, which shows a yellowish-green fluores- 
cence. The property is due to the fact that the wave-lengths of 
part of the invisible, ultra-violet rays of the sunlight are lengthened, 
and a greenish light is therefore in excess. The oxides are U0 2 a 
basic oxide, U 2 3 , U3O8 the most stable oxide, U0 3 uranic anhy- 
dride, and U0 4 a peroxide. 

When the oxide U0 2 is treated with acids, it gives uranous salts 
such as uranous sulphate U(S04) 2 ,4H 2 0. Uranic anhydride and 
uranic acid interact with acids, giving basic salts, such as U0 2 S04, 
3|H 2 0, and U0 2 (N0 3 )2,6H 2 0, which are named uranyl sulphate, 
uranyl nitrate, and so forth. They are yellow in color, with green 
fluorescence. Ammonium sulphide throws down the brown, un- 
stable uranyl sulphide U0 2 S from their solutions. 

Radioactive Elements 

Historical. — We have seen (p. 303) that in an evacuated tube, 
through which an electric discharge is passed, the "rays" emanat- 
ing from the cathode (cathode rays) strike the anti-cathode and 



THE RADIOACTIVE ELEMENTS 



607 



the glass beyond it. They produce in the glass a greenish-yellow, 
fluorescent light. These "rays" were discovered by Sir William 
Crookes (1878), and later were shown to consist of particles of 
negative electricity or electrons, each having a mass about T *Vtf °f 
an atom of hydrogen. Rontgen (1895) accidentally discovered 
that the fluorescent light (X-rays) could penetrate paper, flesh, 
and other materials composed of elements of low atomic weight 
and acted upon photographic plates. In 1896 Henri Becquerel 
observed that minerals containing uranium gave off a sort of 
radiation which could penetrate black paper that was opaque to 
ordinary light and reduce the silver bromide on a photographic 
plate placed beneath the paper. He also discovered that an 
electrometer (Fig. 133), in which the gold leaves had been caused 
to separate by charging with electricity, 
lost its charge rapidly when the uranium 
ore (or salt) was brought near (3-4 cm.) 
to the knob connected with the leaves. 
The uranium material rendered the air 
a conductor ("ionized" the air) and this 
effect permitted the escape of the electric 
charge, which otherwise would have been 
retained for a considerable time. In the 
quantitative measurement of radioactiv- 
ity, we now compare the times required 
for the discharge of an electroscope by different specimens of radio- 
active matter. The presence of 10~ 12 g. of such matter can thus 
be detected. 

The radioactivity of every pure uranium compound is propor- 
tional to its uranium content. The ores are, however, relatively 
four times as active. This fact led M. and Mme. Curie, just after 
1896, to the discovery that the pitchblende residues, from which 
practically all of the uranium had been extracted, were neverthe- 
less quite active. About a ton of the very complex residues 
having been separated laboriously into the components, it was 
found that a large part of the radioactivity remained with the 
sulphate of barium. From this barium sulphate, a product free 
from barium, and at least one million times more active than 
uranium, was finally secured in the form of the bromide. The 
nature of the spectrum and the chemical relations of the element, 




Fig. 133. 



608 COLLEGE CHEMISTRY 

now named radium, placed it with the metals of the alkaline 
earths. The ratio by weight of chlorine to radium in the chloride 
is 35.46 : 113, so that, on the assumption that the element is bi- 
valent, its chloride is RaCl 2 and its atomic weight is 226. With 
this value it occupies a place formerly vacant in the periodic table. 
In 1910 Mme. Curie obtained metallic radium by electrolyzing 
a solution of radium- chloride, using a mercury cathode, and ex- 
pelling the mercury by distillation. It was a white metal (m.-p. 
700°) which, like calcium, quickly tarnished in the air and dis- 
placed hydrogen from water. 

The Nature of the "Rays." — Many properties show that 
the "rays" emitted by compounds of uranium and of radium are 
of three kinds. They are most sharply distinguished from one 
another when allowed to pass through a powerful magnetic field. 
The alpha-rays are positively charged and are bent in one direction 
while the beta-rays are negative and are bent in the other. The 
gamma-rays are not affected. 

The alpha-rays are atoms of helium (p. 336) thrown off in 
straight lines with varying initial velocities, averaging about one- 
tenth that of light (say, 30,000 kilometers per second. The 
a-particles from Ra-C, e.g., 19,220 kilom. per sec). Each such 
atom bears a double positive charge (the unit being the charge on 
a univalent positive ion), and a delicate electroscope readily in- 
dicates the entrance of a single atom. These alpha-particles, 
being each four times as heavy as an atom of hydrogen, plough 
their way through tens of thousands of air-molecules and usually 
go about 3-8 cm. before being stopped. The 
emission of atoms of helium can be detected by 
means of Crookes spintharoscope (Fig. 134). 
The particle of radium bromide is at B, and 



some of the charged helium atoms strike a sur- 
face C covered with zinc sulphide, producing faint flashes of light. 
The lens A magnifies the flashes and the latter can be seen in a 
dark room after the eye has become thoroughly rested (15-20 
minutes). The helium gas given off by radium compounds was 
collected by Soddy working with Ramsay and identified, and its 
rate of production was measured. The amount was equal to 158 
cubic mm. per 1 g. of radium per year. 



THE RADIOACTIVE ELEMENTS 609 

The alpha-particles, in passing through the air-molecules, 
ionize the air, and the ionized ah* has the same power that dust 
possesses (p. 333) of affording nuclei on which moisture can con- 
dense. Hence, when a particle of a radium compound is supported 
in a flask containing air saturated with moisture, and the air is 
suddenly cooled by expansion, the paths of the particles become 
lines of fog. With powerful illumination, the fog-tracks (Fig. 135) 
can be photographed (Wilson), and the lengths of the paths can 
be measured. 

The beta-particles are electrons (p. 303), or unit charges of nega- 
tive electricity, and are shot out with a velocity approaching that 
of light (300,000 kiloms. per sec). They are therefore identical 
with cathode rays, but move many times more rapidly. Being very 
light (weight, T¥ Vo- of an atom of hydrogen), their paths, although 
straight at first, soon become tortuous owing to collisions with 
the relatively massive air-molecules. Half of them are lost after 
going about 4 cm. Their fog tracks are fainter than are those of 
the a-particles and extremely tangled. Being much lighter than 
a-particles, their paths are actually coiled into circles or spirals 
by a magnetic field. 

The gamma-rays are identical with X-rays (vibrations in the 
ether of short wave-length, p. 303), and are presumably produced 
like the latter by the impacts of the electrons on the surrounding 
matter. 

The helium atoms are almost all stopped by a sheet of paper or 
by aluminium foil 0.1 mm. thick. The electrons have greater 
penetrating power, many passing through gold-leaf, but being 
practically all destroyed by a sheet of aluminium 1 cm. thick. 
The gamma-rays (X-rays), however, are able to penetrate rela- 
tively thick layers of metals and other materials of low atomic 
weight. 

One of the most striking facts is that the stoppage by the air 
of so many rapidly moving particles results in the production of 
much heat. One gram of radium would produce about 120 cal. 
per hour. 

Disintegration. — The emission of atoms of helium and of 
electrons was first explained by Rutherford (1902-3), then of 
McGill University, Montreal, as being due to the spontaneous 



610 



COLLEGE CHEMISTRY 




FOG-TRACKS FROM RADIUM (C. T. R. WILSON) 
t, 2. Paths of helium atoms. 3. Part of 2, enlarged. 4. Paths of electrons. 

Fig. 135. 



THE RADIOACTIVE ELEMENTS 611 

disintegration of the atoms of uranium, radium, and other radio- 
active elements. Thus, Rutherford was the first to show that 
radium compounds produced a gaseous substance called the 
radium emanation (niton), which was the residue left after the 
emission of one atom of helium from an atom of radium. This 
gas was itself radioactive and underwent further disintegration, 
depositing a solid radioactive residue on bodies in contact with 
it. Furthermore, every known uranium ore contains radium 
(McCoy) and radium emanation (Boltwood) in amounts propor- 
tional to the uranium content. Also, after the radium has been 
removed, the pure uranium compound gives off at first only 
a-particles, but gradually recovers its whole radioactivity and 
is then found to contain radium emanation once more (Soddy). 
It thus appears that uranium is the starting point, and that the 
disintegration proceeds by steps, producing a number of different 
products. Each of these is formed from one such product and by 
disintegration furnishes another. 

Unlike ordinary chemical change, the rate of disintegration is 
not affected by conditions. It can neither be started nor stopped 
at will. It is no more vigorous at 2000° than at —200°. Other 
changes occur between atoms, these within each atom. 

The law, due also to Rutherford, describing the rate at which 
any one radioactive element disintegrates is simple. Only a 
certain fraction of the whole of any one specimen undergoes the 
change in unit time. Thus, as the total amount diminishes be- 
cause of the change, the amount changing during the next unit 
of time, being a constant fraction of the whole, must be less. 
Hence an infinite time would be required for the complete disin- 
tegration of any one specimen. For convenience, therefore, it is 
sometimes the custom to give as a specific property of each radio- 
active element the time required for the decay of half its amount 
and therefore the loss of half of its radioactivity. More usually, 
the property given is the one called the average life of the element. 
The value of this is equal to the inverse of the fraction disinte- 
grating per unit time, and is about 1.44 times the period of half 
change. Numerically it is the sum of the separate periods of 
future existence of all the atoms divided by the number of such 
atoms present at the starting point. 

Radium emits helium atoms at the rate of 3.4 X 10 10 per gram 



612 COLLEGE CHEMISTRY 

per second. From this fact, we can calculate its average life to 
be about 2400 years. Hence, if it were not continuously being 
produced (from uranium), the whole supply would have been 
exhausted long before the earth reached a habitable condition. 

The Uranium Group of Radioactive Elements. — The fol- 
lowing shows the various elements produced from uranium by 
successive disintegrations. When a helium atom or an electron is 
expelled, the fact is shown by the symbols He and e, respectively. 
The first number below each element is the average life of that 
member of the series (y = year, d = day, h = hour, m = minute, 
s = second). The second number is the atomic weight, obtained 
by subtracting from the at. wt. of uranium (238.2) the weight (4) 
of each helium atom emitted. 

Ui -^He+U-Xj ^€+U-X 2 -+e +U 2 ->He+Ionium 

8X109 y. 35.5 d. 1.65 m. 3X10 6 y. 2X10* y. 

238.2 234.2 234.2 234.2 230.2 

->He+Ra ^He+Niton -^He+Ra-A -+He+Ra-B 

2440 y. 5.55 d. 4.3 m. 38.5 m. 

226 222 • 218 214 

->e+Ra-C^€ +Ra-d ->He+Ra-D ->e +Ra-E 

28.1m. 10~«s. 24 y. 7.2 d. 

214 214 210 210 

->€ +Ra-F->He+Pb(end) 

196 d. 

210 206 

A purified salt of uranium recovers half its activity in about 
three weeks, and reaches full equilibrium in from six months to 
a year. An equilibrium is attained when the speed at which each 
disintegration product is being formed is balanced by the equal 
speed with which it is passing into the next member of the series. 
The complex operations required for studying all the members 
of the series cannot be given here. It may be said, however, 
that a pure uranium salt in solution gives with ammonium car- 
bonate a precipitate which is wholly soluble in excess of the re- 
agent. After about a year, another portion of the same specimen 
leaves a slight precipitate which is insoluble in excess and contains 
the products of disintegration, chiefly U-X which was first obtained 
in this way by Crookes. 



THE RADIOACTIVE ELEMENTS 613 

The radium emanation was shown by Ramsay to be one of the 
inert gases (p. 337), and was renamed niton. Its density was 
determined experimentally with a small sample, using a micro- 
balance capable of weighing to 1/500,000 mgm., and found to be 
222.4 (density of oxygen = 32). 

The end-product of the disintegration is lead, and all uranium 
ores contain lead. Lead from other sources gives a chloride 
PbCl 2 in which 207.20 parts of lead are combined with 2 X 35.46 
parts of chlorine. The atomic weight 207.2 cannot, however, be 
reached by subtracting a whole number of atomic weights of 
helium from the atomic weight of uranium, the number 206 being 
obtained instead. Recently, lead chloride prepared from the lead 
found in various ores of uranium has been analyzed by Richards 
of Harvard, as well as, independently, by two other chemists, 
and the atomic weight of this lead was found to be from 206.4 to 
206.8 in different samples. This lead chloride has properties 
identical with those of ordinary lead chloride and is, therefore, by 
definition, the same substance. Hence these investigations have 
revealed the first known exception to the law of definite pro- 
portions. 

Since the initial (U) and final (Pb) materials are both electri- 
cally neutral, it must be assumed that at some stages more than 
one electron per atom is expelled. 8He++ are lost and therefore 
16e". 

Additional Data, — The yield of radium is very small. 6000 kg. 
of pitchblende, after extraction of the uranium, yield about 
2000 kg. of residue. This affords about 6 to 8 kg. of the mixture 
of radium and barium sulphates, from which 0.2 g. of pure radium 
bromide can be prepared. 

One gram of uranium, after it has produced the equilibrium 
proportion of radium (about 3.2 X 10~ 7 g.), gives off helium at 
the rate of 1 c.c. in sixteen million years. Since the mineral 
fergusonite contains 26 c.c. of accumulated helium for every 
gram of uranium, the samples of this mineral must be at least 
416 million years old. 

The complete disintegration of 1 c.c. of niton to lead would 
deliver about seven million calories, but, of course, the liberation 
of the heat would be spread over a great length of time. 



614 COLLEGE CHEMISTRY 

Chemical Actions of the " Rays." — The radiations which 
are most active in ionizing air and in acting upon photographic 
plates are the a-particles. These particles also cause the flashes 
of light when they encounter zinc sulphide. The radiations 
change the colors of minerals, including gems, and give a deep 
violet color to the glass tube containing the specimen. They 
also turn atmospheric oxygen in part into ozone and, in solution, 
produce traces of hydrogen peroxide in the water. 

The radiations also destroy minute organisms and kill the cells 
of the skin, producing sores. They have been employed in the 
treatment of lupus and of superficial cancerous growths. 

Other Radioactive Series. — Thorium, found as phosphate 
in monazite sand, is also radioactive and furnishes a series of 
disintegration products. The final material is a salt of lead. 
Analysis of the chloride of lead made from traces of the element 
found in all thorium minerals shows that the atomic weight (Soddy) 
is 208.4, while that of ordinary lead is 207.2. The atom of thorium 
(at. wt. 232.4) thus loses 6He (= 6 X 4 = 24) during the disin- 
tegration. There are thus three chlorides of lead with identical 
properties, but different compositions, namely the common one 
207.2 : 2 X 35.46, that from radium 206 : 2 X 35.46, and that 
from thorium 208.4 : 2 X 35.46. 

Actinium and polonium are also radioactive elements, which 
have not yet been fully investigated. The former appears to be 
formed by a second, parallel, disintegration of Ui, and the latter 
in a similar way from Ra-E. Compounds of potassium and 
rubidium show traces of radioactivity. 

Significance of Radioactivity. — The Brownian movement 
(p. 416) has revealed to us bodies intermediate between ordinary 
particles and single molecules, and has enabled us to estimate 
the actual weight of molecules. Radioactivity enables us to 
count charged molecules of helium as they enter the electroscope 
or produce flashes of light on zinc sulphide, and the fog-tracks 
permit us to follow their movements. There is thus now no 
question that molecules and atoms are real. Furthermore, we 
infer that all kinds of atoms are composed of a positive nucleus 
(p. 304) surrounded by electrons, although only the atoms of 



THE RADIOACTIVE ELEMENTS 615 

radioactive elements are unstable. The diameter of the positive 
nucleus of a hydrogen atom is calculated to be about T -g^ of 
that of an electron. Rutherford has confirmed this by actual 
measurement. The atom is thus no longer regarded as being 
solid and continuous in structure. It is mainly a vacuum, con- 
taining a few relatively very minute bodies possessing weight. The 
fact that a-particles are thus able to plough their way through 
molecules of oxygen and nitrogen, being diverted from a straight 
path only when they happen to pass very close to the positive 
nucleus (which, of course, repels the positive a-particles), is no 
longer mysterious. 

Another interesting conclusion has been reached from the ob- 
servation that niton is found in the soil and in many natural 
waters. Calculation shows that the heat given off by the disin- 
tegration of the amounts of radioactive matter known to exist 
in the crust of the earth is alone sufficient to account for the 
maintenance of the temperature of the planet. A globe of the 
size and material of the earth, possessing originally only heat 
energy, and cooling from a white hot condition to the temperature 
of interstellar space, would have passed through the stage of 
habitable temperatures in a much shorter time than that which 
a study of the geological deposits (and the fossils they contain) 
show to have been actually available. The discovery of the 
enormous, but gradually released disintegration energy of the 
radioactive elements enables us now to explain the prolonged 
period during which life has existed on the earth. 

Exercises, — 1. Construct equations, showing the interactions of: 
(a) chromic oxide and aluminium, (6) strontium nitrate and potas- 
sium dichromate in solution, (c) potassium hydroxide and chromic 
hydroxide, and the reversal on boiling, (d) chlorine and potassium 
chromite in excess of alkali (what is the actual oxidizing agent?). 

2. What volume of oxygen at 0° and 760 mm., (a) .is obtain- 
able from one formula-weight of potassium dichromate (par. 4, 
p. 599), (6) is required to oxidize one formula-weight of chromous 
chloride? 

3. To what classes of actions should you assign the three methods 
of making chromic oxide (p. 603)? 



616 COLLEGE CHEMISTRY 

4. Make equations for all the reactions involved in the prepa- 
ration of sodium-diuranate from pitchblende. 

5. How many candle power will be obtained from 50-watt 
carbon and tungsten filament lamps, respectively? 

6. Point out the resemblance, and the differences between the 
reactions of, (a) gold with aqua regia, (b) calcium oxalate with 
hydrochloric acid, (c) barium chromate with nitric acid (p. 601). 



CHAPTER XLIV 

MANGANESE 

The Chemical Relations of the Element. — Manganese 
stands, at present, alone on the left side of the eighth column of the 
periodic table. The right side is occupied by the halogens. It is 
never univalent, as are the halogens, but its heptoxide Mn 2 7 and 
the corresponding acid, permanganic acid HMnC>4, are in many 
ways closely related to the heptoxide of chlorine and perchloric 
acid HCIO4. Of the lower oxides of manganese, MnO is basic, 
and Mn 2 03 feebly basic. Mn0 2 is feebly acidic, Mn03 more 
strongly so, and permanganic acid (from Mn 2 7 ) is a very active 
acid. Contrary to the habit of feebly acidic and feebly basic 
oxides, such as those of zinc, aluminium, and tin, the basic oxides 
of manganese are not at all acidic, and the acidic oxides, with the 
exception of Mn0 2 , are not also basic. There are thus the five 
following, rather well-defined sets of compounds, showing five 
different valences of the element. Of these the first, fourth, and 
fifth are the most stable and the most important. 

1. Manganous compounds, MnO, Mn(OH) 2 , MnS04, etc. 
These compounds resemble those of the magnesium family (and 
those of Fe ++ ) . The salts of weak acids, such as the carbonate and 
sulphide, are easily made, and there is little hydrolysis of the 
halides. The salts are pale-pink in color. 

2. Manganic compounds, Mn 2 3 , Mn(OH) 3 , Mn 2 (S0 4 ) 3 , [MnCl 3 ]. 
The salts resemble the chromic and aluminium salts in behavior, 
but are even less stable than those of quadrivalent lead. They are 
completely hydrolyzed by little water. The salts are violet in 
color. 

3. Manganites, Mn0 2 , H 2 Mn0 3 , CaMn0 3 . The alkali manga- 
nites are strongly hydrolyzed, like the plumbates and the stannates. 

4. Manganates, Mn0 3 , H 2 Mn0 4 , K 2 Mn0 4 . The salts resemble 
the sulphates and chromates, but are much more easily hydrolyzed. 
The free acid resembles chloric acid (p. 314) in that, when it de- 

617 



618 COLLEGE CHEMISTRY 

composes, it yields a higher acid (HM11O4) and a lower oxide 
(Mn0 2 ) . The salts are green in color. 

5. Permanganates, Mn 2 7 , HMn0 4 (hydrated), KMn0 4 . The 
salts resemble the perchlorates, and are not hydrolyzed by water. 
They are reddish-purple in color. 

It will be seen that the element manganese changes its character 
totally with change in valence, and in each form of combination 
resembles some set of elements of valence identical with that which 
it has itself assumed. Since the valence represents the number of 
electrons gained or lost by each atom (p. 322), it is thus evident 
that the chemical properties of an element depend more upon 
the electrical constitution of its atom than upon the atomic weight. 
The latter is a secondary property, dependent on the former 
(c/.p.304). 

Occurrence: the Metal, — The chief ore is the dioxide, pyro- 
lusite Mn0 2 , which always contains compounds of iron. Other 
manganese minerals are: braunite Mn 2 3 ; the hydrated form, 
manganite MnO(OH); hausmannite Mn 3 04; and manganese spar 
MnC0 3 . The metal is most easily made by reducing one of the 
oxides with aluminium by Goldschmidt's method. 

The metal manganese (m.-p. 1260) has a grayish luster faintly 
tinged with red. It is oxidized superficially by air, and easily dis- 
places hydrogen from dilute acids, giving manganous salts. Its 
alloys with iron, such as spiegel iron (5-15 per cent Mn) and ferro- 
manganese (70-80 per cent Mn), are made by using manganese ores 
with the charge in the blast furnace, and are added to the iron in 
making special steels. Manganese steel (7-20 per cent Mn) is 
exceedingly hard, even when cooled slowly. It is used for the 
jaws of rock crushing machinery and for burglar-proof safes. 
Wire made of an alloy called manganin (Cu 84 per cent, Ni 4 per 
cent, Mn 12 per cent), invented by Weston, is used in instruments 
for making electrical measurements, because its resistance does 
not alter with moderate changes in temperature. 

Oxides. — Manganous oxide MnO is a green powder, made by 
reducing any of the other oxides with hydrogen. Hausmannite 
Mn 3 4 is dull red. An oxide having this composition is formed 
when any of the other oxides is heated in air, oxidation or reduction, 



MANGANESE 619 

as the case may be, taking place (cf. p. 575). Manganic oxide 
Mn 2 3 is brownish-black, and is formed by heating any of the 
oxides in oxygen. 

Manganese dioxide Mn0 2 is black, and is most easily prepared 
in pure condition by gentle ignition of manganous nitrate. The 
hydrated forms of the oxide are produced by precipitation, as by 
adding a hypochlorite or hypobromite to manganous hydroxide 
suspended in water. Manganese dioxide is not a peroxide in the 
restricted sense (cf. p. 223). That is to say, it does not contain the 
radical O2 11 and, therefore, does not give hydrogen peroxide. Its 
reaction formula is Mn(0) 2 not Mn(0 2 ) and in double decomposi- 
tions it yields only water H 2 (0). It is used for manufacturing 
chlorine, although electrolytic processes are now driving it out of 
this field. In glass-making (q.v.), it is employed to oxidize the 
green ferrous silicate, derived from impurities in the sand, to the 
pale-yellow ferric compound. The amethyst color of the manganic 
silicate which is formed tends also to neutralize this yellow. It is 
mixed with black paints as a " dryer" (oxidizing agent). 

Manganese trioxide MnOs is a red, unstable powder. Manganese 
heptoxide Mn 2 7 is a brownish-green, volatile oil (see below). 

When any of these oxides is heated with an acid, a manganous 
salt is obtained. Salts of this class are, in fact, the only stable sub- 
stances in which manganese is combined with an acid radical. In 
this action the oxides containing more oxygen than does MnO give 
off oxygen, or oxidize the acid (cf. p. 157). When the oxides are 
heated with bases, in the presence of air, manganates are always 
formed. In this case, with oxides containing a smaller proportion 
of oxygen than Mn0 3 , oxygen is taken from the air. 

Manganous Compounds. — The manganous salts are formed 
by the action of acids upon the carbonate or any of the oxides. 
Thus the chloride MnCl 2 ,4H 2 is obtained in pale-pink crystals 
from a solution made by treating the dioxide with hydrochloric acid 
and driving off the chlorine liberated by oxidation (p. 158). The 
hydroxide Mn(OH) 2 is formed as a white precipitate when a soluble 
base is added to a solution of a manganous salt. This body passes 
into solution when ammonium salts are added, and cannot be 
precipitated in their presence on account of the formation of 
molecular ammonium hydroxide and the suppression of hydroxide- 



620 COLLEGE CHEMISTRY 

ion (cf. magnesium hydroxide, p. 525). The hydroxide quickly 
darkens when exposed to the air and passes over into hydrated 
manganic oxide MnO(OH). 

Manganous sulphate gives pink crystals of a hydrate. Below 6° 
the solution deposits MnS0 4 ,7H 2 0, which is a vitriol (p. 529). 
Between 7° and 20° the product is MnS0 4 ,5H 2 0, asymmetric and 
resembling CuS0 4 ,5H 2 0. Above 25° monosymmetric prisms of 
MnS0 4 ,4H 2 are obtained. These hydrates have different aqueous 
tensions and may be formed from one another by lowering or raising 
the pressure of water vapor around the substance (p. 96). 

Manganous carbonate MnC0 3 is a white powder formed by pre- 
cipitation. The sulphide MnS is obtained as a green, crystalline 
powder by leading hydrogen sulphide over any of the oxides. A 
flesh-colored, amorphous manganous sulphide MnS (often some- 
what hydrated) is more familiar and is precipitated by ammonium 
sulphide from manganous salts. It interacts with mineral acids 
and even with acetic acid, so that it cannot be precipitated by the 
action of hydrogen sulphide on salts (cf. p. 530). When rubbed in 
a mortar it becomes crystalline, and is then green. 

The manganous salts of weak acids, such as the carbonate and 
sulphide, darken when exposed to air and are oxidized, with forma- 
tion of hydrated manganic oxide. As we have seen, manganous 
hydroxide is similarly oxidized and these salts are precisely the ones 
which should furnish the hydroxide by hydrolysis. While there is 
a general resemblance between the manganous salts and the stan- 
nous, chromous, and ferrous salts, the manganous salts of active 
acids are not oxidized by the air as are the corresponding salts of 
the other three metals. 

Manganic Compounds. — The base of this set of compounds, 
manganic hydroxide Mn(OH) 3 , is slowly deposited by the action of 
the air on an ammoniacal solution of a manganous salt in salts of 
ammonium. Manganic chloride MnCl3 is present in the liquid ob- 
tained by the action of hydrochloric acid upon manganese dioxide 
(cf. p. 158), but loses chlorine very readily. 

Manganites. — Although manganese dioxide interacts when 
fused with potassium hydroxide, simple salts derived from 
H 2 Mn0 3 (= H 2 0,Mn0 2 ) or H 4 Mn0 4 ( = 2H 2 0,Mn0 2 ) are not 



MANGANESE 621 

formed. The products are complex, as K 2 Mn 5 0n. Some less 
complex manganites are formed by mixing manganous chloride 
solution with slaked lime, and blowing air through the mass of 
calcium and manganous hydroxides which is thus obtained. Man- 
ganites of calcium, such as CaMn0 3 ( = CaO,Mn0 2 ) and CaMn 2 C>5 
(= CaO,2Mn0 2 ) are thus formed: 

Ca(OH) 2 + 2Mn(OH) 2 + 2 -> CaMn 2 5 + 3H 2 0. 

Manganates. — When one of the oxides of manganese is fused 
with potassium carbonate and potassium nitrate, a green mass is 
obtained. The green aqueous extract deposits potassium man- 
ganate K 2 Mn0 4 in rhombic crystals, which are of the same form as 
those of potassium sulphate, and are almost black: 

K 2 C0 3 + Mn0 2 + O -> K 2 Mn0 4 + C0 2 . 

The acid H 2 Mn0 4 is itself unknown. The potassium salt remains 
unchanged in solution only in presence of free alkali. When the 
concentration of the hydroxide-ion is reduced by dilution, or, better 
still, when a weak acid such as carbonic acid or acetic acid is used 
to neutralize it, the salt is decomposed according to the following 
equation: 

3K 2 Mn0 4 + 2H 2 -> 4KOH + 2KMn0 4 + Mn0 2 . 

That is, a precipitate of manganese dioxide and a solution of 
potassium permanganate are obtained. To make the equation 
(pp. 322-324), we note that in K 2 Mn0 4 we have 2K+ and 40= and 
therefore Mn+tt to secure electrical neutrality. The latter becomes 
MnW" and Mntt Arithmetically 3Mnffi will give 2Mnttt f 
and lMntt. Hence, 3K 2 Mn0 4 are required, and 2KMn0 4 and 
lMn0 2 produced. In terms of the ions the equation is simpler: 

3Mn0 4 = + 2H+ -> 20H~ + 2Mn0 4 _ + Mn0 2 . 

Permanganates. — Potassium permanganate KMn0 4 is made 
by decomposition of the manganate as shown above, and is ob- 
tained, in purple crystals with a greenish luster, by evaporation of 
the solution. To avoid the loss of manganese thrown down as 
dioxide, the action is carried out commercially by passing ozone 



622 COLLEGE CHEMISTRY 

through the solution of the manganate: 2K 2 Mn0 4 + 3 + 
H 2 -» 2KMn0 4 + 2 + 2KOH. Sodium permanganate NaMn0 4 
is made in a similar manner. It is not obtainable in solid form, but 
its solution is known as "Condy's disinfecting fluid." This liquid 
owes its properties to the oxidizing power of the salt. Perman- 
ganic acid is a very active acid, that is, it is highly ionized in 
aqueous solution. A solid hydrate of the acid may be secured in 
reddish-brown crystals by adding sulphuric acid to a solution of 
barium permanganate and allowing the filtrate to evaporate: 
Ba(Mn0 4 ) 2 + H 2 S0 4 + zH 2 <± BaS0 4 j + 2HMn0 4 ,zH 2 0. 

This hydrate decomposes, on being warmed to 32°, and yields 
oxygen and manganese dioxide. When a very little dry, powdered 
potassium permanganate is moistened with concentrated sulphuric 
acid, brownish-green, oily drops of permanganic anhydride (man- 
ganese heptoxide) Mn 2 7 are formed. This compound is volatile, 
giving a violet vapor, and is apt to decompose explosively into 
oxygen and manganese dioxide. Its oxidizing power is such that 
combustibles like paper, ether, and Uluminating gas are set on fire 
by contact with it. 

Potassium Permanganate as an Oxidizing Agent. — The 

actions are different according as the substance is employed (1) in 
acid, or (2) in neutral solution. 

1. In presence of an acid, and an oxidizable body, a manganous 
salt is always formed. The schematic equation, Mn 2 7 — »2MnO+ 
50, shows that every two molecules of the permanganate yield 50 
for oxidizing purposes. Thus, when sulphuric acid is added to 
potassium permanganate solution, and sulphur dioxide is led 
through the mixture, we have: 

2KMn0 4 + 3H 2 S0 4 -* K 2 S0 4 +2MnS0 4 +3H 2 0(+50) (1) 

(50) + 5H 2 S0 3 ->5H 2 S0 4 (2) 

2KMn0 4 +3H 2 S0 4 + 5H 2 S0 3 -> K 2 S0 4 + 2MnS0 4 + 3H 2 + 5H 2 S0 4 

In this case, since sulphuric acid is a product, the preliminary addi- 
tion of the acid was superfluous. In other cases, the partial equa- 
tion (1), showing the available 50, remains the same, while the 
other partial equation varies with the substance being oxidized. 
Thus, with hydrogen sulphide as reducing agent, we have : 

(0) + H 2 S -+ H 2 + S X 5 (20 



MANGANESE 623 

and with ferrous sulphate, we get ferric sulphate: 

2FeS0 4 + H 2 S0 4 (+ 0) -> Fe 2 (S0 4 ) 3 + H 2 X 5 (2") 

As before (2') and (2") must be multiplied throughout by five, 
before summation is made (see also p. 225). 

The quantity of a ferrous salt, or of hydrogen peroxide (p. 225) 
in a sample of a solution may be measured by titrating (p. 257) the 
solution with a standard solution of potassium permanganate until 
the color ceases to be destroyed, and then noting the volume used. 
For iron, the standard solution may be prepared so that 1 cc. will 
oxidize 0.01 g. of Fe ++ . 

2. When dry potassium permanganate is heated, it decomposes 
as follows: 

2KMn0 4 -> K 2 Mn0 4 + Mn0 2 + 2 . 

The neutral solution oxidizes substances which are reducing agents. 
The fingers are stained brown by an aqueous solution, receiving a 
deposit of manganese dioxide, in consequence of the reducing 
power of the unstable organic substances in the skin. The de- 
struction of minute organisms by Condy's fluid results from a 
similar action. When the powdered salt is moistened with glycer- 
ine, the mass presently bursts into flame. 

Analytical Reactions of Manganese Compounds, — The 

ions commonly encountered are manganous-ion Mn++, which is 
very pale-pink in color, permanganate-ion Mn0 4 ~, which is purple, 
and manganate-ion Mn0 4 = , which is green. The manganous 
compounds give with ammonium sulphide the flesh-colored sul- 
phide which is soluble in acids. Bases give the white hydroxide, 
which darkens by oxidation, and is soluble in salts of ammonium. 
All compounds of manganese confer upon the borax bead an 
amethyst color (manganic borate), which, in the reducing flame, 
disappears (manganous borate). A bead of sodium carbonate 
and niter becomes green on account of the formation of the 
manganate. 

Exercises. — 1. What do we mean by saying that, (a) chromous 
chloride is stable (p. 93), but easily oxidized by the air, (6) per- 



624 COLLEGE CHEMISTRY 

manganic acid is an active oxidizing agent in presence of an acid 
(p. 622). 

2. Formulate the oxidations of hydrogen sulphide, of ferrous 
sulphate, of oxalic acid (to carbon dioxide), and of nitrous acid (to 
nitric acid) by potassium permanganate in acid solution. In doing 
so, employ the several methods suggested on pp. 322-326. 



CHAPTER XLV 

IRON, COBALT, NICKEL 

The elements iron (Fe, at. wt. 55.84), cobalt (Co, at. wt. 59), and 
nickel (Ni, at. wt. 58.7) are not corresponding members of succes- 
sive periods, like the families hitherto considered. They are 
neighboring members of the first long period, lying between its first 
and second octaves. 

Iron Fe 

Chemical Relations of the Element. — The oxides and 
hydroxides FeO and Fe(OH) 2 , Fe 2 3 and Fe(OH) 3 are basic, the 
former more strongly so than the latter. The ferrous salts, de- 
rived from Fe(OH) 2 , resemble those of the magnesium group and 
those of Cr 44 " and Mn 44 , and are little hydrolyzed. The ferric 
salts, derived from Fe(OH) 3 , resemble those of Cr 444- and Al 444 * 
and are hydrolyzed to a considerable extent. Iron gives also a 
few ferrates K 2 Fe0 4 , CaFe0 4 , etc., derived from an acid H 2 Fe0 4 
which, like manganic acid H 2 Mn0 4 (p. 621), is too unstable to be 
isolated. Complex anions containing iron, such as the anion of 
K4.Fe(CN) 6 , are familiar, but complex cations containing ammonia 
are unknown. 

Occurrence. — Free iron is found in minute particles in some 
basalts, and many meteorites are composed of it. Meteoric iron 
can be distinguished from specimens of terrestrial origin by the 
fact that it contains 3-8 per cent of nickel. The chief ores of iron 
are the oxides, haematite Fe 2 3 and magnetite Fe 3 4 , and the car- 
bonate FeC0 3 , siderite. The first is reddish and radiated in 
structure; but black, shining, rhombohedral crystals, known as 
specularite, are also found. Hydrated forms, like brown iron ore 
2Fe 2 3 ,3H 2 0, are also common. Siderite is pale-brown in color 
and rhombohedral, like calcite. When mixed with clay it forms 
iron-stone, from which most of the iron in Great Britain, but less 
than one per cent of that in the United States is obtained. Pyrite 

625 



626 



COLLEGE CHEMISTRY 



FeS 2 consists of golden-yellow, shining cubes or pentagonal dodec- 
ahedra. It is used, on account of its sulphur, in the manufacture 
of sulphuric acid, but, from the oxidized residue, iron of sufficient 
purity is obtained with difficulty. Compounds of iron are con- 
tained in chlorophyll and in the blood (haemoglobin), and doubtless 
play an important part in connection with the vital functions of 
these substances. Ammonium sulphide blackens the skin, form- 
ing ferrous sulphide by interaction with organic compounds of iron 
present in the tissues. 

Pure Iron. — Pure iron is obtained by reducing pure ferrous 
oxalate in a stream of hydrogen at a high temperature. It is also 
made by electrolysis of ferrous sulphate solution at 100° between 
iron electrodes. It is silver-white and melts at 1510°. The purest 
iron does not rust in pure cold water, but the impurities in ordinary 
iron act as contact agents and rusting proceeds. 



Metallurgy. — The ores of iron are first roasted in order to 
decompose carbonates and oxidize sulphides, if these salts are 
, present. Coke is then used to reduce the oxides. 

Coal is unsuitable because so much heat is wasted 
in driving out the volatile matter and moisture, 
which are absent from coke. Ores containing lime 
or magnesia are mixed with an acid flux, such as 
sand or clay-slate, in order that a fusible slag may 
be formed. Conversely, ores containing silica and 
clay are mixed with limestone. With proper 
adjustment of the ingredients the process can be 
carried on continuously in a blast furnace (Fig. 
136), an iron structure 40 to 100 feet high, lined 
with firebrick. The solid materials thrown in at 
the top are converted, as they slowly descend, 
completely into gases which escape and liquids 
(iron and slag) which are tapped off at the bot- 
tom. Heated air is blown in at the bottom through 
tuyeres, and the top is closed by a cone which descends for a moment 
when an addition is made to the charge. The gases, which contain 
much carbon monoxide, are led off and used to heat the blast or to 
drive gas-engines. 




Fig. 136. 



iron 627 

The main action takes place between the carbon monoxide, 
present in consequence of the excess of carbon, and the oxide of 
iron: 

Fe 3 4 + 4CO <=± 3Fe + 4C0 2 . 

Since the action is a reversible one, a large excess of carbon mon- 
oxide is required. At 650°, equilibrium is reached with CO : CO2 : : 
1 vol. : If vols., and in practice the proportion of carbon monoxide 
used is from twice to fifteen times as great. Almost 5 tons of air, 
heated in advance to 800°, are required for each ton of iron 
produced. The moisture in this air acts upon the coke, giving 
water-gas (p. 386). This action uses up fuel, and also lowers the 
temperature at the point where it should be highest. In the most 
modern furnaces, therefore, the air-blast is regulated, in accord- 
ance with the amount of moisture present in the air at the time. 
This illustrates the commercial value of even a single improve- 
ment in a chemical operation. If this process were used with 
every blast furnace, an immense sum would be saved, for in the 
United States alone 30 million tons of iron are annually produced 
(1913). This is considerably over 40 per cent of the world's pro- 
duction, 20 per cent being supplied by Germany and 15 per cent 
by Great Britain. 

In the upper part of the furnace, the heat (400°) loosens the 
texture of the ore. Further down, the temperature is higher 
(500-900°), and the carbon monoxide reduces the oxide of iron to 
particles of soft iron. A temperature high enough to melt pure 
iron is barely reached anywhere in the furnace, but, a little lower 
down, by solution of carbon in the iron, the more fusible cast iron 
(m.-p. about 1200°) is formed and falls in drops to the bottom. 
It is in this region also that the slag, essentially a glass (p. 493), 
is produced. If the flux had begun sooner to interact with the 
unreduced ore, iron would have been lost by the formation of the 
silicate. The iron collects below the slag, and the latter flows 
continuously from a small hole. The former is tapped off at 
intervals of six hours or so from a lower opening. As a rule, the 
iron never cools until it has been converted into rails or structural 
iron. In some cases, it is made into "pigs" in a casting machine. 

Cast Iron and Wrought Iron. — Pure iron is not manu- 
factured, and indeed would be too soft for most purposes. Piano- 



628 COLLEGE CHEMISTRY 

wire, however, is about 99.7 per cent pure. The product obtained 
from the blast furnace contains 92-94 per cent of iron along with 
2.6-4.3 per cent of carbon, often nearly as much silicon, varying 
proportions of manganese, and some phosphorus and sulphur. 
The last four ingredients are liberated from combination with 
oxygen by the carbon in the hottest part of the furnace and com- 
bine or alloy themselves with the iron. Cast iron does not soften 
before melting, as does the purer wrought iron (m.-p. 1510°), but 
melts sharply at 1150-1250° according to the amount of foreign 
material it contains. When suddenly cooled it gives chilled cast 
iron which is very brittle and looks homogeneous to the eye, all the 
carbon being present in the form of carbide of iron Fe 3 C (cementite) 
in solid solution in the metal. This solid solution is exceedingly 
hard, but very brittle. By slower cooling, time is permitted for 
the separation of part of the carbon as graphite, which appears in 
tiny black scales, and gray cast iron results. This mixture is much 
softer, on account of the amount of free, relatively pure iron which 
it contains. 

Cast iron is used in making cooking ranges, stoves, pipes, and 
radiators. It expands in solidifying, and so fills every detail of the 
mold. 

Wrought iron, invented by Henry Cort (1784), is made by heat- 
ing the broken pigs of cast iron upon a layer of material containing 
oxide of iron and hammer-slag (basic silicate of iron) spread on the 
bed of a reverberatory furnace (Fig. 116, p. 460). The carbon, 
silicon, and phosphorus combine with the oxygen of the oxide, and 
the last two pass into the slag. The sulphur is found in the slag 
as ferrous sulphide. On account of the effervescence due to the 
escape of carbon monoxide, the process is called "pig-boiling." 
The iron is stirred with iron rods ("puddled") and stiffens as it 
becomes purer, until finally it can be withdrawn in balls ("blooms") 
and partially freed from slag by rolling. The resulting bars are 
repeatedly cut, piled in a bundle, reheated, and rolled. The iron 
now softens sufficiently for welding below 1000° and melts at 
1505°. Its fibrous structure is due partly to the films of slag which 
have not been completely pressed out by the rolling. On account 
of its toughness, wrought iron is used for anchors, chains, and bolts, 
and for drawing into wire. On account of its relative purity 
(99.8-99.9 per cent), it is less fusible than cast iron and is used for 



IKON 



629 



fire bars. The above operations are now largely performed by 
machinery, but have been largely displaced by the Bessemer and 
open hearth processes in which iron of equal purity can be obtained. 

Properties of Steel. — This is a variety of iron almost free 
from phosphorus, sulphur, and silicon. Tool-steel contains 0.9- 
1.5 per cent of carbon, structural steel only 0.2-0.6 per cent, and 
mild steel 0.2 per cent or even less. Steel combines the properties 
of cast and of wrought iron, being hard and elastic, and at the same 
time available for forging and welding when the proportion of 
carbon is not too high. Steel can be tempered (see below). It 
has also a greater tensile strength * than has wrought iron, and it 
can be permanently magnetized. 

Bessemer Process. — Steel is made largely by the Bessemer 
process (Kelly 1852, Bessemer 1855). The molten cast iron is 
poured into a converter (Fig. 137) and a blast of air (a) is blown 
through it. The oxidation of the 
manganese, carbon, silicon, and 
a little of the iron gives out suffi- 
cient heat to raise the temperature 
of the mass above the melting- 
point of wrought iron. The re- 
quired proportion of carbon is 
then introduced by adding pure 
cast iron, spiegel iron, or coke, and 
the contents, first the slag, and 
then the molten steel, are finally poured out by turning the con- 
verter. When the cast iron contains much phosphorus, the oxide 
of this element is reduced again by the iron as fast as it is formed 
by the blast. In such cases a basic lining containing lime and 
magnesia takes the place of the sand and clay lining of the ordinary 
Bessemer converter, and a slag containing a basic phosphate of 
calcium is produced. This modification constitutes what is 
known as the basic or Thomas-Gilchrist process. The slag 
("Thomas-slag") when pulverized forms a valuable fertilizer 

* Tensile strength or tenacity is measured by the weight (in kilos) required to 
break a wire of the metal 1 sq. mm. in section. Lead 2.6, copper 51, iron 71, 
steel 91. 




Fig. 137. 



630 



COLLEGE CHEMISTRY 



(qf.p.488). 
preferred. 



In the United States, the basic open-hearth process is 



Open-Hearth (Siemens- Martin) Process, — In this process 
the cast iron is melted in a saucer-shaped depression (Fig. 138), 
which is lined with sand in the acid process and with lime and 
magnesia in the basic process. Scraps of iron plate (for dilution) 
and haematite, or some other oxide ore, are then added in proper 
proportions. The materials (50-75 tons in one charge) are 
heated with gas fuel for 8-10 hours. To secure economically the 



1 i ' i ' i ' i ' i ' i ' i ' i ' i ' i ' i * i ' i ' i r i ' i ' i ' i ' i T i ' i ' i ' i ' i ' i ' i ' i ' i ' i . ' i ' i H 




Fig. 138. 

high temperature required to keep the product (almost pure iron) 
fused, Siemens devised the method of preheating the fuel gas and 
air by a regenerative device. The spent air and gas pass down 
through a checkerwork of brick. When this becomes heated, the 
valves are reversed, the gas and air now enter through the heated 
brickwork and, after meeting and burning over the iron, pass out 
through the checkerwork on the opposite side, raising its tempera- 
ture in turn. 

The changes are similar to those in the Bessemer process. 
During casting, some aluminium is added to combine with oxygen 
(present as CO) and give sounder ingots. Recently, iron con- 
taining 10-15 per cent of titanium has been added instead. The 



IRON 631 

titanium combines with both nitrogen and oxygen and the com- 
pounds pass into the slag, just as does aluminium oxide. Rails 
made of steel purified with this element are less liable to breakage 
(the commonest cause of wrecks) and are 40 per cent more durable, 
than are ordinary open-hearth rails. 

The advantage of the open-hearth process over that of Bessemer 
is that it is not hurried, and is therefore under better control. The 
material can be tested by sample at intervals until the required 
composition has been reached. The product is of more uniform 
quality. When fine steel is required, electric heating (e.g., in the 
Heroult furnace) permits even more deliberate treatment. 

Bessemer and open-hearth steel is used for heavy and light 
machinery castings and for shafts. It is rolled into rails, and into 
bridge and structural iron. 

Crucible Steel. — For special purposes steel is made in cru- 
cibles of clay (or graphite and clay) in melts of 60-100 pounds. 
" Melting bar," a very pure open-hearth steel, is melted with 
charcoal or with pure pig iron. This steel is employed in making 
razors (1.5 per cent C), tools (1 per cent C), dies (0.75 per cent C), 
pens, needles, and cutlery. 

Tempering. — The carbon in steel (and cast iron) is in the 
form of carbon or of carbide of iron FeaC (6.6 per cent C), dissolved 
in the iron. When white hot steel (up to 2 per cent C) is suddenly 
chilled, there is no time for any changes to occur during the cooling, 
and a solid solution is obtained which is very hard and brittle. 
When, however, the cooling is slow, some of the carbon separates 
in minute crystals of cementite FesC until, at about 700°, there 
remains only about 0.9 per cent carbon in solid solution. At this 
temperature, if sufficient time is allowed, the solid solution sepa- 
rates into a mixture of pure iron (87 per cent) which is soft and 
carbide of iron (13 per cent) which is hard. Steel, when slowly 
cooled, is thus a mixture, and not homogeneous. If, therefore, 
hard chilled steel is heated once more for the purpose of tempering, 
the extent to which the softer material is formed depends upon the 
temperature reached and upon the rate and the duration of the 
cooling process. By varying these, the degree of hardness allowed 
to remain can be adjusted. 



632 COLLEGE CHEMISTRY 

Steel Alloys. — As we have seen, substances such as aluminium, 
titanium, and ferrosilicon are added to iron for the purpose of 
purifying it, and pass in combination into the slag. There are, 
however, regular alloys containing the foreign metal along with 
the iron. Thus, manganese steel (7-20 per cent Mn), made by 
adding spiegel iron or ferromanganese (p. 618) to steel, remains 
hard even when cooled slowly and is used for the jaws of rock- 
crushers and for safes. Chromium-vanadium steel (1 per cent Cr, 
0.15 per cent Va) has great tensile strength, can be bent double 
while cold, and offers great resistance to changes of stress and to 
torsion. It is used for frames and axles of automobiles and for 
connecting rods. Tungsten steel has already been described (p. 
606). Nickel steel (2-4 per cent Ni) resists corrosion, has a high 
limit of elasticity and great hardness, and is used for armor-plate, 
wire cables, and propeller shafts. Invar (36 per cent Ni) is 
practically non-expansive when heated within narrow limits and 
is used for meter-scales and pendulum rods. 

Chemical Properties of Iron. — Although the purest iron 
does not rust in cold water (p. 626), ordinary iron rusts in moist 
air or under water. It probably rusts in water free from carbon 
dioxide, displacing the hydrogen-ion, but the action is greatly 
hastened by the presence of carbonic acid. Rust is a brittle, porous, 
loosely adherent coating of variable composition, consisting mainly 
of a hydrated ferric oxide 3Fe 2 03,H 2 0, which does not protect the 
metal below. Oil protects iron from rusting because, although 
oxygen is more soluble in most oils than in water, and so reaches the 
iron freely, water is not soluble in oil and so moisture is excluded. 

Iron burns in oxygen and it interacts with superheated steam, 
in both cases giving Fe 3 4 . A superficial layer of this oxide ad- 
heres firmly and protects the iron from the action of the air (Barff's 
process iron, or Russia iron). 

Iron displaces hydrogen easily from dilute acids. Steel and cast 
iron, which contain iron, its carbide, and graphite, give with cold 
dilute acids almost pure hydrogen, and the carbide and graphite 
remain unattacked. More concentrated acids, however, particu- 
larly when warm, generate, along with hydrogen, hydrocarbons 
formed by interaction with the carbide (p. 441). The odor of the 
gas is due to compounds of sulphur and phosphorus. 



iron 633 

Although iron acts vigorously on dilute or concentrated nitric 
acid, it is indifferent to fuming nitric acid (N0 2 in solution, p. 348). 
It becomes passive. In this state, it no longer displaces hydrogen 
from dilute acids. If dipped in cupric sulphate solution, it does 
not receive the usual red coating of metallic copper. However, if 
scratched or struck, the passive condition is destroyed, and copper 
begins to be deposited at the point touched and the action spreads 
quickly over the whole surface. No satisfactory explanation of 
this phenomenon has been obtained, although it is shown also by 
chromium, cobalt, and other metals. 

Ferrous Compounds. — Ferrous chloride is obtained as a pale- 
blue hydrate FeCl 2 ,4H 2 (turning green in the air) by interaction 
of hydrochloric acid with the metal or the carbonate. The an- 
hydrous salt sublimes in colorless crystals when hydrogen chloride 
is led over the heated metal. In solution the salt is oxidized by 
the air to a basic ferric chloride : 

4Fe++ + 2 + 2H 2 -> 4Fe+++ + 40H". 

In presence of excess of the acid, normal ferric choride is formed. 
With nitric acid, ferric chloride and nitric oxide are produced (p. 
350). 

Ferrous hydroxide Fe(OH) 2 is thrown down as a white precipitate, 
but rapidly becomes dirty-green and finally brown, by oxidation. 
It dissolves in solutions of salts of ammonium, being like magne- 
sium hydroxide (p. 525), sufficiently soluble in water to require 
an appreciable concentration of OH - for its precipitation. The 
NH4 + from the salts combines with the OH~ formed by the ferrous 
hydroxide to give molecular ammonium hydroxide. Ferrous 
oxide FeO is black, and is formed by heating ferrous oxalate in 
absence of air. It is made also by cautious reduction of ferric 
oxide by hydrogen (at about 300°), but is easily reduced further 
to the metal. It catches fire spontaneously when exposed to 
the air. 

Ferrous carbonate FeC03 is found in nature as siderite, and may 
be made in slightly hydrolyzed form by precipitation. The pre- 
cipitate is white but rapidly darkens and finally becomes brown, 
the ferrous hydroxide produced by hydrolysis being oxidized to the 
ferric condition. The salt interacts with water containing car- 



634 COLLEGE CHEMISTRY 

bonic acid, after the manner of calcium carbonate (p. 383), giving 
FeH 2 (C0 3 ) 2 , and hence is found in solution in natural (chalybeate) 
waters. 

Ferrous sulphide FeS may be formed as a black, metallic-looking 
mass by heating together the free elements. It is produced by 
precipitation with ammonium sulphide, but not with hydrogen sul- 
phide. It interacts readily with dilute acids. The precipitated 
form is slowly oxidized to ferrous sulphate by the air. 

Ferrous sulphate is obtained by allowing pyrites to oxidize in the 
air and leaching the residue: 

2FeS 2 + 70 2 + 2H 2 -» 2FeS0 4 + 2H 2 S0 4 . 

The liquor is treated with scrap iron and the neutral solution evapo- 
rated until a hydrate FeS04,7H 2 0, green vitriol, or " copperas," is 
deposited. The crystals are efflorescent, and become also brown 
from oxidation to a basic ferric sulphate: 

4FeS0 4 + 2 + 2H 2 ->4Fe(OH)S0 4 . 

With excess of sulphuric acid and air, or an oxidizing agent such as 
nitric acid, ferric sulphate is formed. The ferrous sulphate is used 
in dyeing and in making writing-ink. The extract of nut-galls con- 
tains tannic acid, HC14H9O9, which, with ferrous sulphate, gives 
ferrous tannate, a soluble, almost colorless salt. A solution of this 
salt containing gum-arabic and some blue or black dye constitutes 
the ink. When the writing is exposed to the air, the ferrous 
tannate is oxidized to the ferric condition, and the ferric compound 
is a fine, black precipitate (cf. p. 516). The dye is added to make 
the writing visible from the first. Ferrous sulphate is also used in 
the purification of water (p. 560). 

Ferric Compounds. — By leading chlorine into a solution of 
ferrous chloride, and evaporating until the proper proportion 
of water alone remains, a yellow, deliquescent hexahydrate of 
ferric chloride, FeCl3,6H 2 is obtained. When this is heated still 
further, hydrolysis takes place and the oxide remains. When 
chlorine is passed over heated iron, anhydrous ferric chloride 
sublimes in dark green scales, which are red by transmitted light. 
In solution, the salt ? like other ferric salts, can be reduced to the 



iron 635 

ferrous condition by boiling with iron. The same reduction is 
effected by hydrogen sulphide : 

2Fe+++ + Fe->3Fe++. 
2Fe+++ + S=->2Fe++ + S j. 

The ferric ion is almost colorless, the yellow-brown color of solu- 
tions of ferric chloride being due to the presence of ferric hydroxide 
produced by rrydrolysis. The color deepens when the solution is 
heated (increased hydrolysis), and fades again very slowly, by 
reversal of the action, when the cold solution is allowed to stand. 

Ferric hydroxide Fe(OH) 3 appears as a brown precipitate when 
a base is added to a ferric salt. It does not interact with excess 
of the alkali. In this form the substance dries to the oxide with- 
out giving definite intermediate hydrated oxides. The hydrates, 
Fe 2 3 ,2Fe(OH) 3 (brown iron ore) and Fe 2 3 ,4Fe(OH) 3 (bog iron 
ore), however, are found in nature (see Rust, p. 632). The hy- 
droxide passes easily into colloidal solution in a solution of ferric 
chloride, and by subsequent dialysis through a piece of parchment 
the salt can be separated, and a pure colloidal suspension of the 
hydroxide obtained. This suspension, known as dialyzed iron, 
is red in color, shows no depression in the freezing-point, and is 
not an electrolyte. The hydroxide is a positive colloid and is 
coagulated (brown precipitate) by the addition of salts, bivalent 
negative ions being more effective than univalent ones (p. 417). 

Ferric oxide, Fe 2 3 , is sold as " rouge" and " Venetian red." It 
is made from the ferrous sulphate, obtained in cleaning iron ware 
which is to be tinned or galvanized, and in other ways in the arts. 
The salt is allowed to oxidize, and the ferric hydroxide, thrown 
down by the addition of lime, is calcined. The product varies 
in tint from a bright yellowish-red to a dark violet-brown according 
to the fineness of the powder. The best rouge is obtained by 
calcining ferrous oxalate FeC 2 04. This oxide is not distinctly 
acidic, but by fusion with more basic oxides, compounds like 
franklinite Zn(Fe0 2 ) 2 may be formed. It is reduced by hydrogen, 
at about 300° to ferrous oxide, and at 700-800° to metallic iron. 

Magnetic oxide of iron Fe 3 4 or lodestone is found in nature, and 
is formed by the action of air (hammer-scale), steam, or carbon 
dioxide on iron. It forms octahedral crystals, and is a ferrous- 
ferric oxide FeO,Fe 2 3 or Fe(Fe0 2 ) 2 , related to franklinite. 



636 COLLEGE CHEMISTRY 

Ferric sulphide Fe 2 S3 may be made by fusing together the free 
elements, and is obtained also as a precipitate by the addition 
of ammonium sulphide to ferric chloride solution (Stokes). With 
hydrogen sulphide, only sulphur is thrown down (p. 635). 

Ferric sulphate Fe 2 (S0 4 )3 is formed by oxidation of ferrous sul- 
phate, and is obtained as a white mass by evaporation. It gives 
alums (p. 558), such as ferric-ammonium alum (NH 4 ) 2 S0 4 ,Fe 2 (S0 4 )3, 
24H 2 0, which are almost colorless when pure, but usually have a 
pale reddish-violet tinge. 

Pyrite. — The mineral pyrite FeS 2 (Fools' gold) is the sulphide 
of iron which is most stable in the air. It is found in nature in 
the form of glittering, golden-yellow cubes, octahedrons, and pen- 
tagonal dodecahedrons. It is not attacked by dilute acids, but 
concentrated hydrochloric acid slowly converts it into ferrous 
chloride and sulphur. It is reduced by hydrogen to ferrous 
sulphide. 

Cyanides. — When potassium cyanide is added to solutions of 
ferrous or ferric salts, yellowish precipitates are produced, but the 
simple cyanides cannot be obtained in pure form. These precipi- 
tates interact with excess of the cyanide giving soluble complex 
cyanides of the forms 4KCN,Fe(CN) 2 and 3KCN,Fe(CN) 3 . 
These are called ferro- and ferricyanide of potassium, respectively. 
Ferrocyanide of potassium K4Fe(CN) 6 ,3H 2 0, "yellow prussiate 
of potash," is made by heating nitrogenous animal refuse, such as 
blood, with iron filings and potassium carbonate. The resulting 
mass contains potassium cyanide and ferrous sulphide, and when 
it is treated with warm water these interact and produce the ferro- 
cyanide: 

2KCN + FeS -> Fe(CN) 2 + K 2 S, 
4KCN + Fe(CN) 2 -^K 4 .Fe(CN) 6 . 

The salt is made also from the cyanogen contained in crude illumi- 
nating gas. The trihydrate forms large, yellow, monosymmetric 
tables. The solution contains almost exclusively the ions K + and 
Fe(CN) 6 — = , and gives none of the reactions of the ferrous ion Fe++. 
The corresponding acid H 4 .Fe(CN)6 may be obtained as white 
crystalline scales by addition of an acid and of ether (in which the 
substance is less soluble than in water) to the salt, The acid is a 



iron 637 

fairly active one, but is unstable and decomposes in a complex 
manner. Other ferrocyanides may be made by precipitation. 
That of copper Cu 2 .Fe(CN) 6 is brown, and ferric ferrocyanide 
Fe 4 [Fe(CN) 6 ]3 has a brilliant blue color (Prussian blue). The fer- 
rous compound (insoluble) Fe 2 Fe(CN) 6 , or perhaps K 2 FeFe(CN) 6 , 
is white but quickly becomes blue by oxidation. The soluble 
ferrocyanides are not poisonous. 

Ferricyanide of potassium K 3 Fe(CN)6 is easily made from the 
ferrocyanide by oxidation : 

2I^Fe(CN) 6 + C1 2 ->2KC1 + 2K 3 .Fe(CN) 6 , 
or 2Fe(CN) 6 == + Cl 2 -> 2Fe(CN) 6 =- + 2CP. 

It forms red monosymmetric prisms. The free acid H 3 Fe(CN) 6 
is unstable. Other salts may be prepared by precipitation. 
Ferrous ferricyanide Fe 3 [Fe(CN 6 )] 2 is deep-blue in color (Turn- 
bull's blue). With ferric salts only a brown solution is obtained. 
Prussian blue and Turnbull's blue are used in making laundry 
blueing. They are insoluble, but give colloidal suspensions and 
are adsorbed by the material of the cloth. 

Blue-Prints, — Some ferric salts, when exposed to light, are 
reduced to the ferrous condition. Thus, ferric oxalate, in the 
light, gives ferrous oxalate: 

Fe 2 (C 2 4 ) 3 -» 2FeC 2 4 + 2C0 2 . 

When paper is coated with ferric oxalate solution and dried, 
and an ink drawing on transparent paper is placed over the pre- 
pared surface, sunlight will reduce the iron to the ferrous condi- 
tion, excepting where the ink protects it. When the sheet is then 
dipped in potassium ferricyanide solution (developer), the ferric 
oxalate gives only the brown substance which can be washed out. 
But the deep blue, insoluble ferrous ferricyanide is precipitated 
in the pores of the paper where the light has acted. The drawing 
appears white on a blue background. In ordinary blue-print 
paper, ammonium-ferric citrate takes the place of the oxalate, 
and the ferricyanide has already been applied to the paper before 
drying, so that only exposure and washing remain to be done. 
Dilute sodium hydroxide solution decomposes the ferricyanide, 
and is used for writing (in white) on blue-prints. 



638 COLLEGE CHEMISTRY 

Iron Carbonyls, — When carbon monoxide is led over finely 
divided iron at 40-80°, or under eight atmospheres pressure at 
the ordinary temperature, volatile compounds of the composition 
Fe(CO)4, iron tetracarbonyl, and Fe(CO)s, the pentacarbonyl, are 
formed. When the gaseous mixture is heated more strongly, the 
compounds decompose again, and iron is deposited. Illuminating- 
gas burners frequently receive a deposit of iron from this cause. 

Analytical Reactions of Compounds of Iron. — There are 
two ionic forms of iron, ferrous-ion Fe ++ , which is very pale-green, 
and ferric-ion Fe +++ , which is almost colorless. Ammonium 
sulphide gives with the former black ferrous sulphide, which is 
soluble in dilute acids. The hydroxides are white and brown, 
respectively, and ferrous carbonate is white. With ferric salts, 
which are hydrolyzed (about 5%), carbonates yield the hydroxide 
because they neutralize the free acid and displace the equilibrium. 
With ferrocyanide of potassium, ferrous salts give a white, and 
ferric salts a blue precipitate. With ferricyanide of potassium 
the former gives a deep-blue precipitate, and the latter a brown 
solution. Ferric thiocyanate Fe(CNS) 3 is deep-red (p. 182). With 
borax, iron compounds give a bead which is green (ferrous borate) 
in the reducing flame, and colorless or, with much iron, yellow 
(ferric borate) or even brown when oxidized. 

Cobalt Co 

The Chemical Relations of the Element. — Cobalt forms 
cobaltous and cobaltic oxides and hydroxides CoO and Co(OH) 2 , 
C02O3 and Co (OH) 3 , respectively, which are all basic, the former, 
more so than the latter. The cobaltous salts are little hydrolyzed, 
but the cobaltic salts are largely decomposed by water. The 
latter also liberate readily one-third of the negative radical, after 
the manner of manganic salts, becoming cobaltous. Complex 
cations and anions containing cobalt are very numerous and very 
stable. 

Occurrence and Properties. — Cobalt is found along with 
nickel in smalt ite CoAs 2 and cobaltite CoAsS. The pure metal 
may be made by Goldschmidt's process, or by reducing the oxalate, 
or an oxide, with hydrogen. 



COBALT 639 

The metal is silver-white, with a faint suggestion of pink. It is 
markedly crystalline, less tough than iron, and has no commer- 
cial applications. It displaces hydrogen slowly from dilute acids, 
but interacts readily with nitric acid. 

Cobaltous Compounds. — The chloride CoCl 2 ,6H 2 may be 
made by treating the oxide with hydrochloric acid. It forms red 
prisms, and when partially or completely dehydrated becomes 
deep-blue. Writing made with a diluted solution upon paper is 
almost invisible, but becomes blue when warmed and afterwards 
takes up moisture from the breath, and is once more invisible 
(sympathetic ink). Most cobaltous compounds are red when 
hydrated or in solution (Co 4 ^), and blue when dehydrated. By 
addition of sodium hydroxide to a cobaltous salt, a blue basic 
salt is precipitated. When the mixture is boiled, the pink cobalt- 
ous hydroxide Co(OH) 2 is formed. This becomes brown through 
oxidation by the air. It interacts with ammonium hydroxide, 
giving a soluble ammonio-cobaltous hydroxide, which is quickly 
oxidized by the air to an arnmonio-cobaltic compound (see below). 
It dissolves also in salts of ammonium as does magnesium hy- 
droxide (p. 525). When dehydrated it leaves the black cobaltous 
oxide CoO. Cobaltous sulphate, CoS0 4 ,7H 2 0, and cobaltous ni- 
trate, Co(N0 3 ) 2 ,6H 2 0, are familiar salts. The black cobaltous 
sulphide CoS is precipitated by ammonium sulphide from solu- 
tions of all salts, and even by hydrogen sulphide from the acetate, 
or a solution containing much sodium acetate (c/. p. 484). Once 
it has been formed, it interacts very slowly even with dilute hy- 
drochloric acid, having apparently changed into a less active 
form. A sort of cobalt glass, made by fusing sand, cobalt oxide, 
and potassium nitrate, forms, when powdered, a blue pigment, 
smalt, used in china-painting and by artists. 

Cobaltic Compounds. — By addition of a hypochlorite to a 
solution of a cobaltous salt, cobaltic hydroxide Co(OH) 3 , a black 
powder, is precipitated. Cautious ignition of the nitrate gives 
cobaltic oxide C02O3. Stronger ignition gives the commercial 
oxide, which is a cobalto-cobaltic oxide Co 3 4 . Cobaltic oxide 
dissolves in cold hydrochloric acid, but the solution gives off 
chlorine when warmed. By placing cobaltous sulphate solution 



640 COLLEGE CHEMISTRY 

round the anode of an electrolytic cell, crystals of cobaltic sul- 
phate, Co 2 (S0 4 ) 3 , have been made and cobaltic alums have also 
been prepared (Hugh Marshall). 

Complex Compounds. — Potassium cyanide precipitates from 
cobaltous salts a brownish-white cyanide. This interacts with 
excess of the reagent, giving a solution of potassium cobaltocy- 
anide K4.Co(CN) 6 (cf. p. 636). This compound is easily oxidized 
by chlorine, or even when the solution is boiled in the air, and the 
colorless potassium cobalticyanide is formed: 

4K4Co(CN) 6 + 2H 2 + 2 ->4K 3 .Co(CN) 6 + 4KOH. 

The solution gives none of the reactions of CO+++, and with acids 
the very stable cobalticyanic acid, H 3 Co(CN)3, is liberated. 

When acetic acid and potassium nitrite are added to a cobaltous 
salt, the latter is oxidized by the nitrous acid (liberated by the 
acetic acid) and a white complex salt K 3 .Co(N0 2 ) 6 ,nH 2 ( = 
Co(N0 2 ) 3 ,3KN0 2 ), potassium cobaltinitrite, is thrown down. 

Cobaltic salts give with ammonia complex compounds which are 
many and various. The cations often contain negative groups, 
and are such as Co(NH 3 ) 6 +++, Co(NH 3 ) 5 Cl++ and Co(NH 3 ) 5 N0 2 ++. 
Usually the solutions give none of the reactions of cobaltic ions, 
and often fail likewise to give those of the anion of the original salt. 



Nickel Ni 

The Chemical Relations of the Element. — Nickel forms 
nickelous and nickelic oxides and hydroxides NiO and Ni(OH) 2 , 
Ni 2 3 , and Ni(OH) 3 , but only the former are basic. The nickel- 
ous salts resemble the cobaltous and ferrous salts, but are not 
oxidizable into corresponding nickelic compounds. Since there 
are no nickelic salts, there are here no analogues of the cobalti- 
cyanides or the cobaltinitrites. The complex nickelous salts, 
like the complex cobaltous salts, and unlike the complex cobaltic 
salts, are unstable, and so give some of the reactions of Ni"^. 

Occurrence and Properties. — Nickel occurs free in meteor- 
ites and in niccolite NiAs and nickel glance NiAsS. It is now 
manufactured chiefly from pentlandite [Ni,Cu,Fe]S and other 



NICKEL 641 

minerals found at Sudbury (Ontario), and from garnierite, a 
silicate of nickel and magnesium, found in New Caledonia. In 
the former case, the ore is roasted, smelted, and finally bessem- 
erized. The resulting alloy of copper and nickel is much used 
for sheet-metal work (Monel metal, approx. 1:1). Pure nickel 
is separated from the copper by an electrolytic process (p. 511), 
or by the Monde process (see below). 

The metal is white, with a faint tinge of yellow, is very hard, and 
takes a high polish (m.-p. 1452°). It is used in making alloys, 
such as German silver (copper, zinc, nickel, 2:1:1) and the 
"nickel" used in coinage (copper, nickel, 3:1). Although in 
these alloys the red color of the copper is completely lost, the 
copper is simply dissolved, and not combined. Zinc and copper, 
however, give a compound Cu 2 Zn3. Nickel plating on iron is 
accomplished exactly like silver plating (p. 516). The bath con- 
tains an ammoniacal solution of ammonium-nickel sulphate 
(NH4) 2 S04,NiS0 4 ,6H 2 0, and a plate of nickel forms the anode. 

The metal rusts very slowly in moist air. It displaces hydro- 
gen with difficulty from dilute acids; but interacts with nitric 
acid. 

Compounds of Nickel. — The chloride NiCl2,6H 2 is made by 
treating any of the oxides with hydrochloric acid, and is green 
in color (when anhydrous, brown). The sulphate NiS0 4 ,6H 2 0, 
which crystallizes in green, square prismatic forms at 30-40°, is 
the most familiar salt. Nickelous hydroxide, Ni(OH) 2 , is formed 
as an apple-green precipitate, and when heated leaves the green 
nickelous oxide NiO. It dissolves in ammonium hydroxide, 
giving a complex nickel-ammonia cation. It is soluble also in 
salts of ammonium (c/\ p. 525). By cautious ignition of the 
nitrate, nickelic oxide Ni 2 3 is formed as a black powder. The 
oxides and salts, when heated strongly in oxygen, give the oxide 
N13O4. The last two oxides liberate chlorine when treated with 
hydrochloric acid, and give nickelous chloride. Nickelic hydrox- 
ide Xi(OH) 3 is a black precipitate formed when a hypochlorite 
is added to any salt of nickel. Nickelous sulphide is thrown down 
by ammonium sulphide, and behaves like cobaltous sulphide 
(p. 639). It forms a brown colloidal solution when excess of the 
precipitant is used, and is then deposited very slowly. 



642 COLLEGE CHEMISTRY 

Addition of dimethylglyoxime to an ammoniacal solution of a 
salt of nickel gives a brilliant scarlet precipitate of an acid salt: 

Ni(OH) 2 + 2(HON) 2 C 2 (CH 3 ) 2 -*2H 2 + NiH 2 [C 2 N 2 2 (CH 3 ) 2 ] 2 . 

This reaction is not shown by salts of cobalt, especially if oxidation 
to the cobaltic condition has been permitted by contact with air. 

With potassium cyanide and a salt of nickel the greenish nickel- 
ous cyanide, Ni(CN) 2 , is first precipitated. This dissolves in 
excess of the reagent, and a complex salt K 2 Ni(CN) 4 ,H 2 ( = 
2KCN.Ni(CN) 2 ) may be obtained from the solution. This salt 
is of different composition from the corresponding compounds of 
cobalt and of iron, and is less stable. Thus, with bleaching 
powder, it gives Ni(OH) 3 as a black precipitate. When the solu- 
tion is boiled in the air no oxidation to a complex nickelicyanide 
occurs, and indeed no such salts are known. This fact enables 
the chemist to separate cobalt and nickel, for when the mixed 
cyanides are boiled and then treated with bleaching powder, the 
cobalticyanide is unaffected. With potassium nitrite and acetic 
acid no insoluble compound corresponding to that given by cobalt 
salts is formed by salts of nickel. The only known compound 
which could be formed, 4KN0 2 ,Ni(N0 2 ) 2 , is soluble. This action 
also is used for the purpose of separation. The pink color of 
cobalt salts and the green of nickel salts are complementary colors, 
so that, by using suitable proportions of the two, a colorless mix- 
ture can be produced. 

When finely divided nickel, made by reducing the oxide or 
oxalate with hydrogen at a moderate temperature, is exposed to a 
stream of cold carbon monoxide, nickel carbonyl Ni(CO) 4 is formed. 
This is a vapor and is condensable to a colorless liquid (b.-p. 43° 
and m.-p. —25°). The vapor is poisonous. When heated to 
150-180° it is dissociated and nickel is deposited. Cobalt forms 
no corresponding compound. Commercially, pure nickel is sepa- 
rated from copper (and cobalt) in the Monde process by passing 
carbon monoxide over the pulverized alloy, and subsequently 
heating the gas. 

Analytical Reactions of Compounds of Cobalt and Nickel. 

— The cobalt ion Co ++ is pink, and the nickelous ion Ni ++ green. 
The reactions used in analysis have been described in the preceding 



NICKEL 643 

paragraphs. With borax, cobalt compounds give a blue bead 
(cobaltous borate), and nickel compounds a bead which is brown 
in the oxidizing flame and cloudy, from the presence of gray, 
metallic nickel, when reduced. 

Exercises. — 1. What would be the interactions of calcium car- 
bonate when fused with sand and with clay, respectively? 

2. Make equations representing, (a) the oxidation of ferrous 
chloride by air, (b) the hydrolysis of ferrous carbonate and the 
oxidation of ferrous hydroxide, (c) the oxidation of ferrous sul- 
phate with excess of sulphuric acid Tby hypochlorous acid, (d) 
the formation of ferrous and ferric t annates (p. 634), (e) the re- 
duction of ferric chloride by iron and by hydrogen sulphide, 
respectively, (/) the dry distillation of basic ferric sulphate, (g) 
the formation of ferric ferrocyanide and of ferrous ferricyanide. 

3. Explain the solubility of cobaltous and nickelous hydroxide 
in salts of ammonium. 

4. Construct equations to show the formation, (a) of the in- 
soluble potassium cobaltinitrite (nitric oxide is given off), (6) of 
nickelic hydroxide from nickelous chloride and sodium hypo- 
chlorite. Remembering that the hypochlorite is somewhat hydro- 
lyzed, explain why the precipitation in (6) is complete. 

5. Tabulate in detail the chemical relations of the elements 
cobalt and nickel, with especial reference to showing the resem- 
blances and differences. 



CHAPTER XLVI 
THE PLATINUM METALS 

The rarer elements of MendelejefTs eighth group divide them- 
selves into sets of three each. Just as iron, cobalt, and nickel 
have similar atomic weights and much the same specific gravity 
(7.8-8.8), so ruthenium (Ru, at. wt. 101.7), rhodium (Rh, at. wt. 
103), and palladium (Pd, at. wt. 106.7) have specific gravities 
from 12.26 to 11.5. Similarly osmium (Os, at. wt. 191), iridium 
(Ir, at. wt. 193), and platinum (Pt, at. wt. 195.2) form a triad 
with specific gravities from 22.5 to 21.5. Chemically, ruthenium 
shows the closest resemblance to osmium, and both are allied to 
iron. Similarly, rhodium and iridium, and palladium and plati- 
num are natural pairs. 

The six elements are found alloyed in nuggets and particles 
which are separated from alluvial sand by washing. Platinum 
forms 60-84 per cent of the whole. The chief deposits are in the 
Ural Mountains, smaller amounts being found in California, 
Australia, Borneo, and elsewhere. The components are separated 
by a complex series of chemical operations. 

Ruthenium and Osmium. — These metals are gray like iron, 
while the other four are whiter and more like cobalt and nickel. 
They also resemble iron in being the most infusible members of 
their respective sets. Both melt considerably above 2000°. 
They likewise resemble iron in uniting easily with free oxygen, 
while the other four elements do not. Ruthenium gives Ru0 2 
and even Ru0 4 , although the latter oxide is more easily obtained 
indirectly. Osmium gives Os04, "osmic acid," a white crystal- 
line body melting at 40° and boiling at about 100°. The odor 
and irritating effects of the vapor recall chlorine (Gk. 007*17, odor). 
The substance is not an acid, nor even an acid anhydride. The 
aqueous solution is used in histology, and stains tissues in conse- 
quence of its reduction by organic bodies to metallic osmium. 
It is affected particularly by fat. Osmic acid also hardens the 

644 



THE PLATINUM METALS 645 

material without distorting it. Osmium forms also a yellow, 
crystalline fluoride, OsF 8 (m.-p. 34.5°). It will be observed that 
ruthenium and osmium have a maximum valence of eight. 

Rhodium and Iridium. — These metals are not attacked by 
aqua regia, while the other four are dissolved, more or less slowly. 
They are harder than platinum, and iridium is alloyed with this 
metal for the purpose of increasing its resistance to the action of 
acids. They resemble cobalt in having no acid-forming properties. 
The most familiar compounds of iridium are the complex chlorides 
XalrCle (= 3XCl,IrCl 3 ) and X 2 IrCl 6 (= 2XCl,IrCl 4 ). The solu- 
tions of the latter are red, and the acid, chloro-iridic acid H 2 IrCl6, 
is often found in commercial chloroplatinic acid H 2 PtCl6, and 
confers upon it a deeper color. 

Palladium and Platinum. — Palladium is the only metal of 
this family which is attacked by nitric acid. Palladium and plati- 
num form -ous and -ic compounds of the forms PdX 2 and PdX4, 
respectively. The oxides PdO and PtO and corresponding hy- 
droxides are basic. When quadrivalent, the metals appear chiefly 
in complex compounds, like H 2 .PtCl 6 and H 2 .PdCl 6 , in which the 
metal is in the anion. Platinum gives also platinates derived from 
the oxide Pt0 2 . 

Palladium. — This metal (m.-p. 1549°), named from the planet- 
oid Pallas, is noted chiefly for its great tendency to absorb hy- 
drogen. When finely divided, it takes up about 800 times its 
own volume. The amount absorbed varies continuously with 
the concentration (pressure) of the hydrogen, although not ac- 
cording to a uniform rule, and the product is in part at least a 
solid solution. When a strip of palladium is made the cathode 
of an electrolytic cell, over 900 volumes of hydrogen may be 
occluded. This absorbed hydrogen, in consequence of the cata- 
lytic influence of the metal, reacts more rapidly than does the gas, 
and consequently a strip of hydrogenized palladium will quickly 
precipitate, from solutions of their salts, copper and other metals 
less electropositive than hydrogen and will reduce ferric and other 
reducible salts: 

CuS0 4 + H 2 -> H 2 S0 4 + Cu, or Cu++ + H 2 -* 2H+ + Cu. 
2FeCl 3 + H 2 -> 2FeCl 2 + 2HC1, or 2Fe+++ + H 2 -> 2Fe++ + 2H+. 



646 COLLEGE CHEMISTRY 

Platinum* — This metal (dim. of Sp. plata, silver) is grayish- 
white in color, and is very ductile. At a red heat it can be welded. 
It does not melt in the Bunsen flame, but fuses easily in the oxyhy- 
drogen jet (m.-p. 1755°). On account of its very small chemical 
activity it is used in electrical apparatus and for making wire, 
foil, and crucibles and other vessels for use in laboratories. It 
interacts with fused alkalies, giving platinates. The oxygen acids 
are without action upon it, but on account of the tendency to 
form the extremely stable complex ion PtCl 6 == (p. 520), the free 
chlorine and chloride-ion in aqua regia convert it into chloro- 
platinic acid H 2 PtCl6. 

The metal condenses oxygen upon its surface and it dissolves 
hydrogen. The finely divided forms of the metal, such as platinum 
sponge made by igniting ammonium chloroplatinate (NH^PtCle, 
platinum black made by adding zinc to chloroplatinic acid, and 
platinized asbestos made by dipping asbestos in a solution of chlo- 
roplatinic acid and heating it, show this behavior very conspicu- 
ously. They cause instant explosion of a mixture of oxygen and 
hydrogen, in consequence of the heat developed by the rapid 
union of that part of the gases which is condensed in the metal. 
A heated spiral of fine platinum wire will continue to glow if im- 
mersed in the mixture of methyl alcohol vapor and air (oxygen), 
formed by placing a little of the alcohol in the bottom of a beaker. 
Some cigar-lighters work on this principle. The heat is developed 
by the interaction between the substances, which takes place with 
great speed at the surface of the platinum. Platinum sponge is 
used as a contact agent in making sulphur trioxide (p. 279). 

Platinum was the only otherwise suitable substance which had the 
same coefficient of expansion as glass, and it was consequently fused 
into incandescent bulbs and furnished the electrical connection 
with the filament in the interior. Recently, however, a less ex- 
pensive substitute has been found. Large amounts are also con- 
sumed in photography and by dentists. It is used also in making 
jewelry, and in Russia for coinage. The price of the metal is 
subject to great variations, since a rainy season in the Caucasus 
will render larger amounts accessible to the miners; but, on the 
whole, the many applications which have been found for it have 
quintupled its price in the last thirty years. The price is now 
about twice that of gold. 



THE PLATINUM METALS 647 

When special resistance to chemical or mechanical influences is 
required, as in standard meters for international reference, or 
points of fountain pens, the alloy with iridium is employed. 

Compounds of Platinum. — Platinous chloride is made by 
passing chlorine over finely divided platinum at 240-250", or by 
heating chloroplatinic acid to the same temperature. It is greenish 
and insoluble in water, but forms with hydrochloric acid the 
soluble chloroplatinous acid H 2 PtCl 4 . Potassium chloroplatinite 
K 2 PtCl4 is used in making platinum prints. Bases precipitate 
black platinous hydroxide Pt(OH) 2 , which interacts with acids but 
not with bases. Gentle heating gives the oxide PtO and stronger 
heating the metal. With potassium cyanide and barium cyanide 
soluble platino-cyanides, K 2 Pt(CN) 4 ,3H 2 and BaPt(CN) 4 ,4H 2 0, 
are formed. These substances, when solid, show strong fluores- 
cence (p. 606), converting X-rays as well as ultra-violet rays into 
visible radiations. The barium salt is used to coat screens on 
which the shadows cast by X-rays are received. 

Chloroplatinic acid H 2 PtCl6,6H 2 is made by treating the metal 
with aqua regia, and forms reddish-brown deliquescent crystals. 
With potassium and ammonium salts, it yields the sparingly solu- 
ble, yellow chloroplatinates K 2 PtCl 6 and (NH 4 ) 2 PtCl 6 (cf. p. 452), 
in solutions of which the platinum migrates towards the anode and 
silver salts precipitate Ag 2 PtCl 6 and not silver chloride. Platinic 
chloride PtCL* is made by heating chloroplatinic acid in a stream 
of chlorine at 360°. When dissolved in water, it combines to 
form H^.PtCLiO, with the platinum in the negative ion. Bases 
interact with chloroplatinic acid, giving a yellow or brown pre- 
cipitate of platinic hydroxide Pt(OH) 4 . This substance interacts 
with bases to give platinates, like Na^HioPtaOi^H^O. Both sets 
of platinum compounds interact with hydrogen sulphide, giving 
the sulphides PtS and PtS 2 , respectively. These are black powders 
which dissolve in yellow ammonium sulphide solution, much 
as do the sulphides of gold, arsenic, and other metals, giving am- 
monium sulphoplatinates. 



APPENDIX 
I. The Metric System 

Length. 1 meter (1 m.) = 10 decimeters = 100 centimeters (100 
cm.) = 1000 millimeters (1000 mm.). 

1 kilometer = 1000 meters (1000 m.) = 0.6214 miles. 
1 decimeter = 0.1 m. = 10 centimeters = 3.937 inches. 
1 meter = 1.094 yards = 3.286 ft. = 39.37 in. 

Volume. 1 liter = 1000 cubic centimeters (1000 c.c.) = a 
cube 10 cm. X 10 cm. X 10 cm. 

1 liter = 0.03532 cu. ft. = 61.03 cu. in. = 1.057 quarts (U. S.) 
or 1.136 quarts (Brit.) = 34.1 fl. oz. (U. S.) = 35.3 oz. (Brit.). 
1 fl. ounce (U. S.) = 29.57 c.c. 1 ounce (Brit.) = 28.4 c.c. 
1 cu. ft. = 28.32 liters. 

Weight. 1 gram (g.) = wt. of 1 c.c. of water at 4° C. 1 kilo- 
gram = 1000 g. 

1 gram = 10 decigrams = 100 centigrams (100 cgm.) = 1000 
milligrams (1000 mgm.). 

1 kilogram = 2.205 lbs. avoird. (U. S. and Brit.). 
1 lb. avoird. = 453.6 g. 

1 oz. avoird. (U. S. and Brit.) = 28.35 g. 100 g. = 3.5 oz. 
1 nickel (U. S.) weighs 5 g. 1 halfpenny (Brit.) weighs 
5 to 5.7 g. 

II. Scale of Hardness 

Each of the following minerals will scratch the surface of a 
specimen of any one preceding it in the list. 

1. Talc 6. Felspar 

2. Gypsum (or NaCl) 7. Quartz 

3. Calcite (or Cu) 8. Topaz 

4. Fluorite 9. Corundum 

5. Apatite 10. Diamond 

648 



APPENDIX 



649 



Glass is slightly scratched by 5, and easily by those following. 
Glass will not scratch 5 distinctly, but will scratch those preceding 5. 
A good knife scratches 6 slightly, but not those following. 
A file will scratch 7, but not those following. 



III. Temperatures Centigrade and Fahrenheit 

Upon the centigrade scale, the freezing-point of water is 0° C. 
and the boiling-point 100° C. Upon the Fahrenheit scale, the 
same points are 32° F. and 212° F., respectively. The same inter- 
val is thus 100° on the one scale and 180° on the other. The degree 
Fahrenheit is therefore j§# or | of 1° Centigrade. Any tempera- 
tures can be converted by using the formulae: 

C.° = | (F.° - 32), F.° = | (C.°) + 32. 

The following table (IV) contains the temperatures from 0° C. 
to 35° C, with the corresponding values on the Fahrenheit scale 
(32° F. to95°F.). 



IV. Vapor Pressures of Water 

Both the Fahrenheit (F.) or ordinary and the Centigrade (C.) temperatures are given. 



Temperature. 




Temperature. 








Pressure, mm. 




Pressure, mm. 


F. 


C. 


F. 


C. 


32° 


0° 


4.6 


71.6° 


22° 


19.7 


41 


5 


6.5 


73.4 


23 


20.9 


46.4 


8 


8.0 


75.2 


24 


22.2 


48.2 


9 


8.6 


77.0 


25 


23.6 


50.0 


10 


9.2 


78.8 


26 


25.1 


51.8 


11 


9.8 


80.6 


27 


26.5 


53.6 


12 


10.5 


82.4 


28 


28.1 


55.4 


13 


11.2 


84.2 


29 


29.8 


57.2 


14 


11.9 


86.0 


30 


31.5 


59.0 


15 


12.7 


87.8 


31 


33.4 


60.8 


16 


13.5 


89.6 


32 


35.4 


62.6 


17 


14.4 


91.4 


33 


37.4 


64.4 


18 


15.4 


93.2 


34 


39.6 


66.2 


19 


16.3 


95.0 


35 


41.8 


68.0 


20 


17.4 








69.8 


21 


18.5 


212^6 


ioo 


760.0 



650 



COLLEGE CHEMISTRY 



V. Order of Activity of the Metals 
(Electromotive Series) 



Each metal, when placed in 
metals following it in the list, 
deposits it in the free condition 

For explanation of potential 
see pp. 539-547. 



a solution of a salt of one of the 
displaces the second metal and 
(see pp. 60, 260, 438, 531). 
differences (electromotive series), 



Potassium 


Manganese 


Tin 


Mercury 


Sodium 


Zinc 


Lead 


Silver 


Barium 


Chromium 


Hydrogen 


Palladium 


Strontium 


Cadmium 


Copper 


Platinum 


Calcium 


Iron 


Arsenic 


Gold 


Magnesium 


Cobalt 


Bismuth 




Aluminium 


Nickel 


Antimony 





INDEX 



*** Acids are all listed under " acid " and salts under the positive radical. 



Acetone, 394, 408 
Acetylene, 378, 392, 394, 400 
formula of, 109 
torch, 394 
Acid, acetic, 407, 467 
antimonic, 589 
arsenic, 585 
arsenious, 586 
boracic, 431 
boric, 431 
bromic, 318 
carbolic, 349 
carbonic, 383 
chlorauric, 356 
chloric, 314 

chloroplatinic, 356, 647 
chlorous, 314 
chromic, 597 
disulphuric, 285 
formic, 385, 412 
hydrazoic, 340, 345 
hydriodic, 202 
hydrobromic, 198 
hydrochloric, 141 
hydrochloric, properties, 146 
hydrocyanic, 420 
hydrofluoboric, 431 
hydrofluoric, 206 
hydrofluosilicic, 427 
hydrosulphuric, 269 
hypochlorous, 161, 307, 309 
hyponitrous, 357 
iodic, 318 

metaphosphoric, 368, 371, 417 
metastannic, 569, 571 
nitric, 347 

fuming, 348 

graphic formula, 358 

oxidizing actions, 354 

synthetic, 352 

test, 351 
nitrosylsulphuric, 281 
nitrous, 356 

orthophosphoric, 368, 370 
08mic, 644 



Acid, oxalic, 385, 413 

palmitic, 412 

perchloric, 314, 315 

perchromic, 224 

permanganic, 622 

persulphuric, 291 

phosphoric, 368 

phosphorous, 372 

picric, 349 

pyrophosphoric, 368, 371 

prussic, 420 

selenic, 294 

silicic, 428 

a-stannic, 570 

sulphuric, 279, 280 
graphic formula, 291 
properties, 285 

sulphurous, 288 

tannic, 634 

thiosulphuric, 290 
Acidic oxides, 94 
Acidimetry, 255 
Acids, 52, 94 

and anhydrides, 316, 369 

fractions ionized, 241 

non-ionic formation, 261 

of constant boiling-point, 145 

organic, 412 

properties in solution, 210 
Actinium, 614 
Actions, non-ionic, 260 

reversible, 177 
Activity, acids, 242 

apparent, 180 

bases, 242 

chemical, 38, 172 

of ionogens, 242 

order of, metals, 59 
non-metals, 548 
Adsorption, 408, 419 
Affinity, chemical, 180 
Agate, 427 
Air, a mixture, 333 

components of, 328, 333 

liquid, 334 



651 



652 



INDEX 



Air, water vapor in, 88 

weight of 22.4 L, 101 
Alabaster, 485 
Alcohol, denatured, 407 

ethyl, 406, 407 

methyl, 408 
Alcohols, 413 
Alkalimetry, 255 
Allotropic modifications, 222 
Alloys, 435, 503, 591 

acid-resisting, 596 

anti-friction, 588 
Alum, 82, 558 

chrome, 603 
Aluminates, 557 
Aluminium, 554 

carbide, 391 

compounds, 556 
Aluminothermy, 556 
Alundum, 558 
Amalgam, sodium, 345 
Amalgams, 435 
Amethyst, 427 
Ammonia, 340 

household, 344 

properties, 342 

-soda process, 461 
Ammonio-copper salts, 504, 506, 507, 
509 

-silver salts, 515 
Ammonium amalgam, 455 

carbonate, 211 

compounds, test for, 345 

cyanate, 421 

hydroxide, 344 

molybdate, 605 

nitrate, 357 

nitrite, 338 

salts of, 345, 453 

sulpharsenate, 587 

sulphides, 454 

sulphostannate, 572 

thiocyanate, 421 
Ammono-compounds, 535 
Amorphous bodies, 97 
Ampere, 137 
Amylase, 406 
Analysis, qualitative, 537 

volumetric, 257 
Analytical reactions, aluminium, 565 

ammonium, 455 

arsenic family, 593 

cadmium, 531 

calcium, 495 

calcium family, 498 

*** Acids are all listed under "acid' 



Analytical reactions, chromium, 604 

cobalt and nickel, 642 

copper, 510 

iron, 638 

lead, 580 

magnesium, 526 

manganese, 623 

mercury, 536 

potassium, 453 

silver, 518 

sodium, 465 

tin, 572 

zinc, 530 
Anhydride, and acid, 369 

and acid or salt, 316 

chromic, 598, 601 

permanganic, 622 
Anhydrides, 94 
Anions, 237 
Anode, 237 
Anthracene, 411 
Antimony, 587 

compounds, 588 
Apatite, 362, 486 
Aq, 52 

Aqua regia, 356 
Aqueous tension, 87, 649 

correction for, 73 

hydrates, 96 
Argentic, see silver 
Argon, 335 
Arsenic, 582 

white, 585 
Arsine, 583 
Asphalt, 391 
Assaying, 521 
Atmosphere, 328 
Atom, constitution, 304 
Atomic numbers, 303 

weight of a new element, 118 

weights, 41, 103, inside rear cover 
advantages of, 107 
Atoms, 43 
Attributes, 19 
Avogadro, 77 

B.T.U., 409 

Babbitt's metal, 588 
Baking powders, 463 

soda, 463 
Barium, 496 

peroxide, 222 
Barometer, 71 
Bases, 94 

fractions ionized, 242 
and salts under the positive radical. 



INDEX 



653 



Bases, properties in solution, 211 

Basic oxides, 94 

Batteries, see cells 

Bead tests, 372, 433 

Beer, 406 

Benzene, 392, 411 

Benzine, 391 

Beryllium, 523 

Bessemer process, 629 

Bicarbonates, 3S3 

Birkeland-Eyde process, 353 

Bismuth, 591 

Black-lead, 378 

Blast lamp, 398 

furnace, 626 
Blau gas, 396 
Bleaching, 311 

hydrogen peroxide, 224 

powder, 309, 312, 484 
Blue-prints, 637 
Blue-stone, 95, 509 
Body, definition of, 4 
Boiling-point, acid of constant, 145 

solutions, 216 
Bone black, 408 
Borax, 431, 432, 465 
Bordeaux mixture, 509 
Boron, 430 
Brass, 503 
Bromine, 193 

oxygen acids, 318 

properties, 195 
Bronze, 503, 

Brownian movement, 416 
Bunsen flame, 398 
Burette, 256 
Butter, 414 
By-product coke, 340, 411 

Cadmium, 530 
Caesium, 453 
Calcining, 275 
Calcite, 83 
Calcium, 474 

bicarbonate, 384 

bisulphite, 290, 402 

carbide, 379, 394 

carbonate, 476 

chloride, 475 

cyanamide, 487 

fluoride, 204, 475 

hydroxide, 477 

hydride, 475 

light, 58 

nitrate, 353 



Calcium, oxalate, 478-484 

oxide, 477 

phosphate, 362, 486 

phosphide, 365 

silicate, 493 

sulphate, 84, 485 

sulphide, 486 
Calculations, formulae from data, 45' 

involving weights, 66 

volumes, 115, 149 
Calomel, 533 
Calorie, 85 
Calorimeter, 174 
Camphor, 291 
Caramel, 405 
Carbides, 378, 379 
Carbohydrates, 402 
Carbon, 375 

dioxide, 381 

as plant food, 387 
uses, 385 

disulphide, 379 

gas, 395 

monoxide, 385 

prints, 600 

tetrachloride, 379, 392 
Carbona, 379 
Carbonates, 383 
Carbonyl chloride, 163, 387 
Carborundum, 380 
Carnallite, 194, 445 
Catalysts, definition, 29 

negative, 288 
Cathode, 237 
Cations, 237 

recognition of, 537 
Cause, 173 
Cell, Clark, 548 

combination, 540, 541 

concentration, 541, 550 

displacement, 540, 544, 547 

Edison, 579 

gravity, 547 

oxidation, 541, 545 

potential differences, 546 

storage, 577 

Weston, 548 
Celluloid, 359 
Cellulose, 402 
Cement, 562 
Cerium, 580 

oxide, 397 
Chalk, 476 
Charcoal, 408 

as adsorbent, 419 



*** Acids are all listed under " acid " and salts under the positive radical. 



654 



INDEX 



Chamber process, 281 
Chemical change, complete, 178 

energy and, 167 

reversible, 177 

speed of, 173 

varieties, 7, 14, 55, 147, 148, 
197 

ionic, 259, 270, 504 
Chemical changes, concurrent, 
317 

consecutive, 289 
Chemical equilibrium, 177 

applications, 184 

displacement of, 185, 203 

history, 187 

in ammonia, 343 

temperature and, 188 
Chemical relations, 163, 192 

halogens, 319 

sulphur family, 295 
Chlorates, 313 
Chlorides, preparation, 146 

solubilities, 164 
Chlorine, 154 

dioxide, 314 

monoxide, 307 

not a bleacher, 312 

oxides and acids, 306 

properties, 159 

-water, 161, 310 
Chloroform, 392 
Chromic anhydride, 598, 601 

compounds, 602 
Chromite, 597 
Chromium, 595 
Chromous compounds, 604 
Chromyl chloride, 601 
Clarke, F. W., 17 
Clay, 430, 561 
Coal, 409 

analysis, 409 

calorific power, 409 
Coagulation treatment, 91 
Cobalt, 638 

compounds, 639 
Coke, 411 

by-product ovens, 340 
Colemanite, 431 
Collie, 15 
Collodion, 359 
Colloids, 403, 415-420, 562 

arsenious sulphide, 586 

ferric hydroxide, 635 
Columbium, 594 
Combination, 7 

*** Acids are all listed under ' 



Combining proportions, measure- 
ment, 33 
Combustion, 35 

spontaneous, 37 
Complex salts, 505 
166, Components, 4 

Composition, definition of, 34 
Concentration, chemical equilibrium 
315, and, 181 

gases, 76 
Concurrent reactions, 315, 317 
Conditions, 20 

Conductivity of ionogens, 239 
Congo red, 258, 565 
Consecutive reactions, 289 
Conservation of mass, 18 
Constant, equilibrium, 184 

ion-product, 470 

ionization, 238 

molecular depression, 214 
Constituents, 8 
Contact action, definition, 29 
Copper, 501 

compounds, 504 

pyrites, 264 

refining, 511 
Copperas, 634 
Corn syrup, 404 
Cordite, 359 

Corrosive sublimate, 533 
Coulomb, 237 
Couples, 549 
Cream of tartar, 463 
Critical temperature, 78 
Cryolite, 204, 554 
Crystal structure, 306 
Crystallization, water of, 96 
Crystals, 82 
Cupellation, 513 
Cupric bromide, 249 

chloride, 504 

nitrate, 211 

oxide, 34, 507 

sulphate, 95, 509 
Cuprous chloride, 504 
Cyanogen, 420 

Davy, 15 

Deacon's process, 156, 177 
Decantation, 12 
Decomposition, 14 
Decrepitation, 449 
Deliquescence, 134 
Density, gases, 73 
I relative, of gases, 152 
acid" and salts under the positive radioal. 



INDEX 



655 



Density, solutions, 138 

vapor, 74 
Depilatory, 486 
Dextrin, 404 
Dextrose, 404 
Dewar flask, 335 
Dialysis, 416 
Diamond, 376 
Diffusion, 57, 78 
Digestion, 422, 423 
Dihydrol, 138 

Dimorphous substances, 266 
Displacement, 55 

ionic, 258 
Dissociation, 93 

cases of, 116 

in solution, 210 
Distillation, fractional, 390 
Double decomposition, 147 

ionic formulation of, 251 
Drummond light, 58 
Dust in air, 328 
Dyes, 411, 419, 563 
Dynamite, 359 

Earthenware, 561 
Earths, metals of the, 553 

rare, 304, 553 
Efflorescence, 96 
Electric energy, units, 539 
Electric furnace, 363, 377, 380 
Electric waves, wireless, 303 
Electrolysis, 55, 155 

explanation of, 232 

products of, 227 

quantities of electricity, 237 
Electrolytic refining, 511, 549, 574 
Electromotive chemistry, 539 

series, 260, 547, 650 
Electrons, 322, 609 

and ions, 235 
Electrophoresis, 417 
Electroplating, 510, 516 
Electrotyping, 510, 516 
Element, used in two senses, 16 
Elements, common, 17 

non-metallic, 94, 296 

metallic, 94, 296 
Emeralds, 524 
Emulsion, 121, 418 
Energy, chemical change and, 167 

conservation of, 170 

internal, 172 

source of world's, 388 
Enzymes, 405 

*** Acids are all listed under " acid 



Equations, 44 

concurrent reactions, 317 

partial, 194 

thermochemical, 174 

writing, 50, 52, 97, 115, 276, 322- 
326, 360 
Equilibrium, chemical, 177 
characteristics, 179 

constant, 184 

displacement of, 90, 143 

ionic, 238, 247, 466 
displacement of, 248 
saturated solutions, 469 

liquid and vapor, 89 

saturated solution, 130 

three characteristics of, 89 
Equivalent weights, 65 
Esters, 413 
Ethyl acetate, 413 
Ethylene, 392, 393, 400 
Explanation, 22 
Explosives, 358 

Fats, 414, 423 
Fehling's solution, 404, 507 
Felspar, 3, 430 
Fermentation, 406 
Ferric compounds, 634 

thiocyanate, 182 
Ferrosihcon, 425 
Ferrous compounds, 633 

sulphide and acids, 272 
Ferrovanadium, 594 
Fertilizers, 387, 451, 488 

ammonium sulphate, 340 

nitrogen, 339 
Filter, Pasteur, 92 
Filtration, 12, 91 
Fire-damp, 391 
Fire extinguishers, 379, 384 
Fixation of nitrogen, 352 
Flame, 396 

blast-lamp, 398 

Bunsen, 398, 400 

cone-separator, 401 

luminosity, 399 

structure, 399 
Flotation, froth, 503 
Flour, wheat, 3 
Fluor-spar, 475 
Fluorine, 204 
Fluorite, 475 
Flux, 438 
Foods, 421-423 

fuel value of, 95 
and salts under the positive radical. 



656 



INDEX 



Formula, reaction, 95 
Formulae, 44 

and valence, 63 

from data, 45 

making, 97 

graphic, 291, 358, 372 

molecular, 109 
Formulation, of chemical equilibrium, 
183 

of double decomposition, 251 

of neutralization, 254 

of precipitation, 252 
Fractions ionized, data, 241 
Freezing-point, definition, 86 

of solutions, 134, 213 
Freezing mixtures, 134 
Froth notation, 503 
Fructose, 404 
Fuels, 410 
Furnace, electric, 363, 377, 380, 488 

G.M.V., 101, 103 
Galena, 573 
Gallium, 553 
Galvanized iron, 528 
Garnet, 82 
Gas, blau, 396 

coal, 409 

-lighters, 580 

oil, 396 

perfect, 79 

producer, 385 

water, 386 

carburetted, 395 
Gases, density, 73 

laws of, 70, 76 

liquefiabilities, 278 

liquefaction, 78, 334 

measurement, 70 

mixed, 72 

solubilities of, 278 
Gasoline, 391 
German silver, 504 
Germanium, 567 
Glass, 493 

etching, 206, 494 

quartz, 428 

uranium, 606 

water, 428 
Glauber's salt, 96, 464 
Glucinum, 523 
Glucose, 403 
Gluten, 3 
Glycerine, 413 
Gypsum, 485 

*tf* Acids are all listed under "acid 



Gold, 518 

compounds, 520 
Gram-molecular volume, 102 
Granite, 2, 430 
Grape-sugar, 404 
Graphic formulae, 291 
Graphite, 376, 377 
Guano, 338, 348 
Guncotton, 350, 358 
Gunpowder, 449 
Gypsum, 84, 485 

Halogen family, 192, 207 

chemical relations, 319 
Hardness, scale of, 648 
Heat, animal, 36 

of neutralization, 255 

of solution, 125 

of vaporization, 86 

thermochemistry, 174 
Heavy-spar, 496 
Helium, 15, 336, 608 
Household ammonia, 344 
Humidity, 328 
Hydrates, 95 
Hydrazine, 340, 345 
Hydrocarbons, 389 

cracking of, 395 

unsaturated, 392 
Hydrogen, 49 

chemical properties, 58 

commercial sources, 56 

dissociation of, 113 

history, 49 

-ion, 246 

nascent, 360 

physical properties, 56 

preparation, 49, 51, 53, 55, 56 
Hydrogen bromide, 196 
Hydrogen chloride, properties, 144 

composition by volume, 164 

preparation, 141 
Hydrogen iodide, 201 

peroxide, 222 

sulphide, 267 
and iodine, 202 
Hydrolysis, 197, 437 

of salts, 271 
Hydrolyte, 475 
Hydrone, 50 
Hydroxide-ion, 246 
Hydroxylamine, 340 
Hypo, 464 
Hypochlorites, 308 
Hypochlorous anhydride, 307 
and salts under the positive radical. 



INDEX 



657 



Ice, 85 

heat of fusion, 86 
Indicators, 257 
Indium, 553 
Infusorial earth, 428 
Ink, printers' and India, 398 

writing, 634 
Internal rearrangement, 148, 421 
Invar, 632 
Iodic anhydride, 319 
Iodine, 198 

chlorides of, 208 

union of hvdrogen and, 203 
Iodoform, 392 
Iodothyrene, 199 
Ion-product constant, 470 
Ionic equilibrium, 466 
Ionic substances, 245 

names of, 236 
Ionization, 226 

activity and, 242 

constant, 238 

degree of, 240, 241 

oxidation and, 322 

questions answered, 234 
Ionogens, classes, 245 

non-ionic formation, 260 
Ions, and electrons, 235 

migration of, 229, 231 
Iridium, 645 
Iron, 625 

carbonv 

cast, 627 

chemical properties, 626, 632 

compounds, 633 

galvanized, 528, 550 

metallurgy, 626 

passive, 633 

Russia, 632 

wrought, 628 
Isomers, 421 
Isoprene, 392 

Javel, eau de, 448 

Kaintte, 451 
Kaolin, 430, 561 
Kerosene, 391 
Kindling temperature, 35 
Kipp apparatus, 54 
Krypton, 337 

Lactose, 404 
Lakes, 565 
Lampblack, 398 

*** Acids are all listed under "acid 



Lard, 414 
Laughing gas, 358 
Lavoisier, 6, 15, 26 
Law, 21 

Avogadro's, 77 

Boyle's, 71, 76 

Charles', 72, 76 

combining weights, 42 

conservation of mass, 18 

Dalton's, 72 

definite proportions, 17, 614 

Dulong and Petit's, 108 

Faraday's, 232 

Gay-Lussac's, 98 

Henry's, 128 

Le Chatelier's, 190 

mass action, 182 

multiple proportions, 47 
Law of, chemical change, 7 

component substances, 3 

molecular concentration, 182 

partition, 129 
Law, periodic, 300 

van't Hoff's, 188 
Laws of gases, deviations from, 78, 

79 
Le Blanc process, 460 
Lead, 573 

compounds, 574 

from radium, 613 

from thorium, 614 

pencils, 378 

red, 575 

white, 577 
Lime, 477 

light, 58 
Liquids, associated, 206 

molecular relations, 81 
Litharge, 575 
Lithium, 465 
Litmus, 258 
Lithopone, 497 
Lomonssov, 4, 5, 15, 80 

Magnalium, 555 
Magnesium, 524 

compounds, 525 

nitride, 339 
Malachite, 502 
Maltose, 404 
Manganese, 617 
Manganic compounds, 620 
Manganin, 618 
Manganites, 621 
Manganous compounds, 619 
and salts under the positive radical. 



658 



INDEX 



Marsh gas, 391 
Marsh's test, 584 
Matches, 365 
Matrix, 264 

Matter, structure of, 74 
Mayow, 5, 14, 25 
Melting-point, definition, 86 
Mendelejeff, 298 
Mercuric oxide, 14, 27 
Mercury, 532 

compounds, 533 
Metallic elements, 94 

chemical relations, 436 
Metals, electromotive series of, 260 

extraction, 438 

melting-points, 435 

occurrence, 437 

order of activity, 59, 650 

physical properties, 434 

potential differences, 547 

world's production, 436 
Methane, 378, 391 
Methyl orange, 258 
Methylated spirit, 407 
Metric system, 648 
Mica, 2, 430 
Microcosmic salt, 371 
Migration of ions, 229 
Milk, 422 
Minium, 575 
Mirrors, silvering, 517 
Mixture, 4 
Moisture, surface, 88 
Molar solutions, 125 
Molar weight, 102 
Molasses, 405 
Mole, 102 

number of molecules in, 103 
Molecular equations, interpretations, 

115 
Molecular formulae, 109 
Molecular theory, 74 

gases, 74 

history, 80 

liquids, 81 

liquid and vapor, 88 

of solutions, 125 

solids, 81 
Molecular weights, 100 

by freezing-point, 134 

in solution, 214 

of elements, 110 
Molybdenum, 604 
Monde process, 642 
Monel metal, 641 



Mortar, 478 

Moseley's atomic numbers, 303 

Naphtha, 391 

Naphthalene, 411 

Neon, 15, 337 

Neutralization, formulation of, 254 

heat of, 255 
Nickel, 640 

carbonyl, 642 

sulphate, 83 
Niton, 337, 612 
Nitric anhydride, 349 
Nitric oxide, 350 
Nitrides, 339 
Nitro-lime, 487 
Nitrogen, 338 

iodide, 346 

tetroxide, 351 

trichloride, 346 
Nitroglycerine, 350, 358 
Nitrosyl chloride, 356 
Nitrous anhydride, 357 

oxide, 357 
Nomenclature, 64, 306 
Non-metallic elements, 94 

potential differences, 548 
Normal solutions, 124 

Oil gas, 396 

Oil, cotton seed, 414 

of vitriol, 285 

olive, 414 
Oleum, 285 

Open-hearth process, 630 
Osmium, 644 

Osmotic pressure, 125, 135 
Ostwald, 21 
Oxidation, 36 

always with reduction, 269 

and reduction, 320-326 
Oxides, acidic, 94 

basic, 94 

order of stability, 60 
Oxidizing agents, explanation of ac- 
tivity, 221 
Oxygen, 25 

chemical properties, 31 

history of, 25 

physical properties, 30 

preparation of, 26 

uses of, 37 

why atomic weight 16, 46 
Oxone, 28 
Ozone, 219 



**& Acids are all listed under " acid " and salts under the positive radical. 



INDEX 



659 



Paint, 579 

lithopone, 497 

luminous, 486 

permanent white, 496 
Palladium, 57, 645 
Paper, 402 

sizing, 560 
Paraffin, 390 
Paris green, 508 
Parke's process, 513 
Pasteur filter, 92 
Pauling process, 353 
Pearl ash, 450 
Perchlorates, 315 
Perchloric anhydride, 316 
Periodic system, 297, inside rear cover 
Permutite, 491 
Peroxidates, 223 
Peroxides, 223 
Petrol, 391 
Petrolatum, 391 
Petroleum, 390 

refining of, 393 
Phenolphthalein, 258 
Phosgene, 163 
Phosphate rock, 362 
Phosphine, 365 
Phosphonium iodide, 366 
Phosphorescence, 364 
Phosphoric anhydride, 367 
Phosphorite, 362, 486 
Phosphorus, 362 

acids of, 368 

pentachloride, 117, 367 

pentasulphide, 373 

pentoxide, 367 

trichloride, 367 

tribromide, 197 

tri-iodide, 201 

vapor, 117 
Photography, 517, 600, 637 
Picture restoring, 224 
Plants and carbon dioxide, 387 
Plaster of Paris, 485 
Plastics, 359 
Platinum, 646 

as catalyst, 59 
Plumbago, 375 
Polarization, 548 
Polonium, 614 
Polysulphides, 274 
Porcelain, 561 
Potash, 450 
Potassium, 443 

alum, 558 

*** Acids are all listed under " acid ' 



Potassium, bisulphite, 452 

bitartrate, 463 
Potassium bromide, 445 

sulphuric acid on, 196 
Potassium carbonate, 450 

chlorate, 27, 313, 448, 469 

chloride, 444 

chromate, 597 

cobalticyanide, 640 

cobaltinitrite, 640 

cuprocyanide, 508 

cyanate, 421 

cyanide, 451 

dichromate, 597 

ferricyanide, 637 

ferrocyanide, 451, 636 

fluorides, 445 

hydroxide, 446 

hypochlorite, 308 
Potassium iodide, 445 

sulphuric acid on, 201 
Potassium manganate, 621 

nitrate, 83, 448 

oxides, 447 

perchlorate, 315, 448 

permanganate, 83, 157, 225, 621 

sulphate, 451 

sulphides, 452 

thiocyanate, 421 

tri-iodide, 200 
Potential differences, single, 546, 547 
Potentials, discharging, 548 
Precipitates, description of, 147 
Precipitation, in presence of acids, 484 

formulation of, 252 

theory of, 478 
Pressure, osmotic, 126, 135 

partial, 72 

solution, 128 

vapor, 87 
Priestley, 14, 26 
Problems, arithmetical, 45, 66, 115, 

149 
Producer gas, 385 
Properties, specific chemical, 30 

specific physical, 19, 30, 31 
Proteins, 422, 486 

test, 350 
Prussian blue, 637 
Pyrene, 379 
Pyrite, 264, 275, 636 

Qualitative analysis, 537 
Quantitative experiments, 33, 34, 35 
Quartz, 3, 83, 427 

and salts under the positive radical. 



660 



INDEX 



Quartz glass, 428 
Quicklime, 477 

Radicals, 53, 212 

positive and negative, 53 

valence of, 62 
Radioactive elements, 606 
Radioactivity, significance, 614 

uranium group, 612 
Radium, 608 
Reactions, concurrent, 315, 317 

consecutive, 289 
Reaction formula, 95 
Realgar, 586 

Reduction and oxidation, 320-326 
Refrigeration, 342 
Relations, chemical, 163, 192 
Reversible actions, 93 
Rey, Jean, 5 
Rhodium, 645 
Rittman's process, 391, 395 
Roasting, 275 
Rochelle salts, 463 
Rock crystal, see Quartz 
Root nodules, 339 
Rubber, synthetic, 392 
Rubidium, 453 
Rusting, 1, 4, 5, 6 
Ruthenium, 644 

Salammoniac, 453 
Saleratus, 450 
Salt, common, 82 
Saltpeter, air, 353 

Bengal, 347 

Chile, 347, 448 
Salts, 149 

double, 245 

fractions ionized, 242 

ions of, 246 

non-ionic formation, 261 

mixed, 245 

properties in solution, 210 
Sandstone, 430 
Saponification, 415 
Saponin, 420 
Scheele, 26 
Schoenite, 451 
Schwerin process, 562 
Selenite, 485 
Selenium, 293 
Sewage, 36 

Siemens-Martin process, 630 
Silicates, 430 
Silicon, 425 

*** Acids are all listed under "acid 



Silicon, dioxide, 427 

tetrafluoride, 206, 427 
Silk, imitation, 359, 507 
Silver, 512 

complex compounds, 514 

salts, 515 
Slag, 438 

Smokeless powder, 359 
Soap, 412, 415, 490 

cleansing power, 418 

salting out, 417 
Soda, washing, 96 
Soda-water, 382 
Sodium, 457 

-ammonium phosphate, 371 

bicarbonate, 462 

carbonate, 96, 460 

chloride, 82, 458, 472 

cyanide, 488 

dichromate, 598 

hydride, 458 

hydroxide, 459 

hyposulphite, 464 

iodate, 318 

metaphosphate, 371 

nitrate, 459 

orthophosphates, 370, 464 

oxides, 459 

palmitate, 412 

peroxide, 28, 222 

persulphate, 291 

silicate, 428 

sulphate, 96, 463 
solubilities, 132 

tetraborate, 432, 465 
Sodium thiosulphate, 290, 464 
Solids, molecular relations, 81 
Solubilities, 131, inside front cover 
Solubility, gases, 128 

measurement of, 123 

product, 471 

temperature and, 130 

units to express, 124 
Solution, 121 

as a process, 127 

dissociation in, 210 

freezing-points, 134, 213 

heat of, 125 

insoluble salts by acids, 481 

molecular theory, 125 

physical or chemical, 138 

pressure, 128 

rule for, 479 

saturated, 123, 128 
definition, 133 
and salts under the positive radical. 



INDEX 



661 



Solution, solid, 122 

supersaturated, 132 

volume changes in, 138 
Solutions, boiling-points, 135 

see Colloids 

densities, 138 

freezing-points, 134 

molar, 125 

normal, 124 

standard, 257, 623 

vapor tension, 134 
Solvay process, 461 
Specific heat, 85, 108 
Spintharoscope, 608 
Sponges, 428 
Stability, chemical, 38 

compounds, 93 
Stalactites, 476 
Starch, 3, 403, 423 
States of matter, 86 
Stationary layers, 328, 398 
Steam, 86 
Stearin, 414 
Steel, 629-632 

alloys, 632 

chromium-vanadium, 632 

manganese, 618, 632 

nickel, 632 

tungsten, 606 
Stereotype metal, 511 
Stibine, 588 
Strontium, 495 
Structure of matter, 74 
Sublimation, 199 
Substance, 2 

simple, 15 
Substitution, 162 
Sucrase, 406 
Sucrose, 404 
Sugar, 84 

cane, 404 

invert, 405 
Sugars, 404 
Sulphates, 288 
Sulphate-ion, 287 
Sulphides, 270 

action of acids on, 271 

insoluble, classification of, 273 

solubilities of, 531 
Sulphites, 290 
Sulphur, 264 

properties, 266 

vapor, 117 
Sulphur, acids of, 280 

dioxide, 35, 275, 277 

*** Acids are all listed under "acid 



Sulphur, monochloride, 291 

trioxide, 279 
Sulphuryl chloride, 278, 291 
Superphosphate, 487 
Symbols, 44 

T. N. T., 349 
Tanning, chrome, 600 
Tantalum, 594 
Tartar-emetic, 589 
Tellurium, 294 

Temperature, and speed of reaction, 
59, 187 

conversion table, 649 

critical, 78 
Tempering, 631 
Tensile strength, 629 
Tension, aqueous, 87, 649 
Thallium, 553 
Theory, molecular, 74, 80 
Thermite, 556 
Thermochemistry, 174 
Thorium, 580 

oxide, 397 
Tin, 567 

compounds, 570 

-plate, 550, 569 

see Stannous and Stannic 
Titanium, 580, 630 
Titration, 256 
Toluene, 349, 392 
Transition points, 86 
Trinitrotoluene, 349 
Tungsten, 605 
Turnbull's blue, 637 

Ultramarine, 563 
Ultramicroscope, 416 
Units, electrical, 237 

of measurement, 648 
Uranium, 606 

radioactivity of, 612 
Urea, 421 

Valence, 61 

and formula?, 63 

and oxidation, 321 

definition, 62 

exceptional, 64 

how ascertained, 63 
Vanadium, 593 
Vapor, density, measurement of, 74 

equilibrium with liquid, 89 

pressure, 87 

saturated, 88 

and salts under the positive radical 



662 



INDEX 



Varnish, black, 398 

Vaseline, 391 

Ventilation, 328 

Verdigris, 508 

Vermilion, 535 

Vitriols, 529 

Volt, 237 

Volume, gram-molecular, 102 

Volumetric analysis, 257 

Warnings, 68, 102, 111, 119, 133, 
175 260 

Washing soda, 96, 462 
Water, 85 

as solvent, 90 

chemical properties, 92 

coagulation process, 560 

composition of, 98 

dihydrol, 138 

domestic, purification, 312 
Water gas, 386 

carburetted, 395 

**?* Acids are all listed under " acid ' 



Water glass, 428 
hard, 415, 489-493 
of crystallization, 96 
physical properties, 85 
vapor tensions, 649 

Waters, natural, 91 

Weights, atomic, 103 
equivalent, 65 
molar, 102 
molecular, 100 

Welsbach lamp, 397 
mantles, 580 

Whisky, 406 

Witherite, 496 

Wood, distillation of, 408 

Wood's metal, 591 

X-rays, 303, 609 
Xenon, 337 

Zinc, 527 

compounds, 528 
Zymase, 406 
and salts under the positive radical. 



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INTERNATIONAL ATOMIC WEIGHTS (1917) 



= 16. 



Aluminium 
Antimony . 

Argon 

Arsenic 

Barium 

Bismuth . . . 

Boron 

Bromine . . . 
Cadmium. . 
Caesium..- 
Calcium.. . 

Carbon 

Cerium 

Chlorine . . , 

Chromium j 

Cobalt... 

Columbiun 

Copn 

Dysfl 

Erbiij 

Euro; 

FluoJ 

Gadc 

GaUi 

Gem 

Gluci 

Gold 

Helh; 

Holn 

Hydi 

Indk 

Iodin 

Iridii 

Iron. 

Kr>T 
Lant! 
Lead 
Lithi 
Lute< 
Magi 
Manj 
Merc 



..Al 
..Sb 
..A 
..As 
..Ba 
..Bi 
..B 
..Br 
..Cd 
.Cs 



27.1 

120.2 

39.88 

74.96 

137.37 

208.0 

11.0 

79.92 

112.40 

132.81 



Molybdenum Mo 

Neodymium Nd 

Neon Ne 

Nickel Ni 

Niton (radium emanation) Nt 

Nitrogen N 

Osmium Os 

Oxygen O 

Palladium Pd 

Phosphorus .... m^m P 



SEAL GH^IISTRY FOR COLLEGES 



Smith, Alexander, 



QD 33 .Sm5 




= 16. 

96.0 
144.3 

20.2 

58.68 
222.4 

14.01 
190.9 

16.00 
106.7 

31.04 
195.2 
9.10 
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226.0 
102.9 



